
Class 

Book_ 



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MODERN CHEMISTRY 



WITH ITS PRACTICAL 
APPLICATIONS 



BY 



FREDUS N. PETERS, A.M. 

INSTRUCTOR IN CHEMISTRY IN CENTRAL HIGH SCHOOL, 

KANSAS CITY, MISSOURI 

AUTHOR OF " EXPERIMENTAL CHEMISTRY," ETC. 




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NEW YORK 
MAYNARD, MERRILL, & CO. 

1901 
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THE LIBRARY OFJ 
CONGRESS, 

Two Copies Received 

JUN. 21 1901 

^Copyright entry 

CLASS a XXc. N«. 

COPY 3. 



Copyright, 1901, by 
MAYNARD, MERRILL, & CO. 



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Narfoooti $ress 

J. S. Clashing & Co. — Berwick & Smith 
Norwood Mass. U.S.A. 



PREFACE 

In preparing this book for use in secondary schools, I have 
endeavored to look at the science from the viewpoint of the 
students themselves. The fault with many texts upon the 
same subject is that the position of the learner has been disre- 
garded; the books have been encyclopsedic ; they have pre- 
sented a great number of facts as a skeleton or framework, 
but this skeleton has not been clothed with muscle and ani- 
mated with life. No more fascinating subject finds a place 
in the curricula of our secondary schools, yet to the average 
student chemistry is too often but an irksome task. 

In the present work I have omitted much that is often 
given in an elementary text, while at the same time entering 
more into that detail which gives lively interest to the subject. 
It has been my aim to show, whenever possible, the practical 
application of the science to the everyday affairs of life; in 
other words, to emphasize industrial and commercial chem- 
istry. At the same time the fundamental principles of the 
science have not been forgotten ;. on the contrary, they have 
been emphasized even more than is usual in an elementary 
chemistry. This has been rendered possible by the omission 
of much that can never be either of interest or of value to the 
beginner. 

Eecognizing the fact that science must be taught inductively 
by experiment, some authors have assumed that the student 
must gain all his knowledge of chemistry in this way. No 
greater mistake could be made. The science has been hun- 
dreds of years in reaching its present development, and much 
must be accepted by the student without any effort to work 
it out for himself. In this text a large amount of experi- 

3 




4 PREFACE 

mental work is given, sufficient to meet the requirements of 
all our best colleges. The experiment is always supplemented 
by notes and suggestions which enable the student to draw 
proper conclusions, and give him such information as he 
cannot hope in his limited time to gain for himself. It will 
be noticed also that the laboratory directions are largely in 
the form of questions, so as to compel even the least ener- 
getic students to secure the benefits of personal investigation. 
This plan is always followed except in cases where the student 
would be in danger of going astray. 

To the pedagogical treatment of the difficult parts of the sci- 
ence I desire to call attention. The subject has been presented 
much in the same way as in my own classes, where the method 
has met with success. Beginning with the study of that 
most familiar of substances, Water, the text enters into a 
discussion of its composition and then proceeds to a detailed 
statement of its constituents. This work is prefaced merely 
by a short chapter connecting the science of Chemistry with 
that of Physics and by one chapter upon Valence. This some- 
what difficult subject of Valence is introduced so early in 
the book in order to avoid many difficult questions that must 
arise when it is deferred. Although introduced early in the 
book, Valence is treated in so simple a way and with so many 
graphic illustrations that the student can hardly fail to grasp 
its meaning. 

I recognize the demands, coming from all higher institutions 
of learning, for more quantitative work, and believe I have 
fully met all such requirements. The student of the subject, 
as taught hitherto, has been in danger of coming to the con- 
clusion that very little in chemistry is exact; whereas nothing 
could be further from the truth. The pupil is shown this 
by the series of quantitative experiments which have been 
carefully worked out in the laboratory. For this work I have 
sought to make use of such apparatus only as may be or 
should be found in any secondary school. 



PREFACE 5 

To the regular text is appended detailed instruction for 
various laboratory manipulations, preparing solutions, etc. A 
chapter has also been added for the benefit of any who may 
desire to continue along qualitative lines the work introduced 
at various places in the text. 

I desire to acknowledge, with gratitude, the valuable sug- 
gestions offered by Professor Irving P. Bishop, of the State 
Normal School, Buffalo, and Professor M. D. Sohon, of the 
Boys' High School, New York City, both of whom have read 
the book critically in manuscript ; furthermore, I wish to say 
that to many of my students of the past I am indebted for 
descriptions, original with them, and more appropriate than 
any I have found in any text. To Dr. Paul Schweitzer, for 
many years Professor of Chemistry in the Missouri State Uni- 
versity, the true friend of the student, to whom I owe much 
for his great sympathy and encouragement, and his words of 
fatherly advice, I desire to express especial gratitude. Finally, 
I acknowledge with pleasure the help and inspiration I have 
gained in my private study and research from the writings of 
those who have been permitted to drink long and deep from 
this fountain of science. 



TO THE TEACHER 

It is not expected that everything given in the text will be 
demanded of the pupil, unless possibly in reviews. Some of 
the manufacturing processes, for example, I have deemed of 
sufficient importance and interest to be given; yet it may seem 
best in the judgment of the instructor to omit these. 

The experiments, as a rule, may be performed by the stu- 
dents, and apparatus is suggested which most schools will be 
able to provide. The number of pupils will determine to 
some extent what experiments should be performed by the 
teacher and what by the students themselves. If the classes 
are small, so that the teacher can give very close personal 
attention, the pupils may attempt almost all; on the other 
hand, if the classes are large, it may be necessary for the 
teacher to perform some of the more difficult and dangerous ex- 
periments himself. On pages 355 to 380 will be found many 
useful suggestions to the student for the care and manipula- 
tion of apparatus, making up of solutions, etc. These should 
be read before the student begins his work in the laboratory, 
and frequent reference should afterward be made to them. 

It is presumed that a school year of nine or ten months will 
be given to the work in this text, but by omitting some of the 
less important elements, much of the theory and many of the 
practical applications of chemical science may be obtained in 
five or six months. Besides the various chapters devoted to 
the fundamental laws of chemistry, special study should be 
given to the following elements and a few of their important 
compounds : hydrogen, oxygen, nitrogen, fluorine, chlorine, 
bromine, iodine, carbon, sulphur, sodium, calcium, zinc, lead, 
and iron. 

6 



MODERN CHEMISTRY 

CHAPTER I 

INTRODUCTION 

1. With what is Chemistry concerned ? — Nature pre- 
sents herself in ever-changing forms, and to one who is 
not familiar with these variations she is a mystery. The 
untaught inhabitant of the tropics, who has never been 
beyond the confines of his native state, taken to a colder 
climate would see no relation between the snowflake or 
the icy covering of our northern rivers and the rain-drop 
as it falls upon his native hills. To him they are entirely 
different substances. 

2. So the diamond, the filling of the ordinary " lead " 
pencil, and the coal that we burn in our furnaces seem 
altogether dissimilar, and yet they are practically the same 
thing. Likewise the emery with which the seamstress 
sharpens her needle and the mechanic his tools, and such 
valuable stones as the oriental emerald and the ruby, though 
seemingly so different, have really the same composition. 
The purpose of the science of chemistry is the investiga- 
tion of the objects that lie all about us in nature, the study 
of their composition and of their relations to one another, 
the explanation of the various phenomena in connection 
with them, and the ability to apply this knowledge to 
practical uses. 

3. Importance of the Subject. — A knowledge of chemis- 
try adds a charm to many of the common things of life, 

7 



8 MODERN CHEMISTRY 

clothing them with new beauty. Later it will be noticed 
that the science of chemistry enters into all or nearly all 
of the great manufacturing industries of the world, and 
that without the application of its laws and principles all 
such enterprises would result in failure. Whether studied, 
therefore, for its intellectual or for its practical value, it 
is of the greatest importance. 

MATTER, THEORIES — PHYSICAL AND CHEMICAL CHANGES 

4. Some Theories of Matter. — There have been men in 
.11 ages who believed that all substances .whatsoever might 

be resolved back into one particular kind of matter; that 
)Y subjecting this matter to different conditions an end- 
less variety of modifications would result. * To illustrate : 
here is a bar of steel ; by submitting it to varying processes 
it is made into saws, knives, needles, watch-springs, pens, 
md thousands of other articles. It is claimed that in the 
same manner the one elementary substance, in undergoing 
different treatments by the forces of nature, appears in 
:he endless variety of substances about us. It has never 
jeen possible, however, for those who hold this view to 
prove it by any experiments. Neither does it seem that 
the phenomena of nature require or even admit of any 
such explanation. 

5. Elements. — What seems a more reasonable view, 
and one that has come nearer to demonstration, is that 
matter is composed of a large number of simple substances, 
and that these combined in different ways produce an 
infinite number of substances. According to this view 
there are about seventy-five simple substances, called ele- 
ments, which cannot be, or at least never have been, 
divided into two or more simpler substances. Numerous 
experiments of every character, by means of the strongest 



MOLECULES AND ATOMS 11 

sisting of two elements, sodium and chlorine, always in the 
ratio of 46 to 71 by weight. Likewise, water is a com- 
pound, containing one part of hydrogen to eight of oxygen. 

8. Divisibility of Matter. — Physics teaches us that 
matter is anything that occupies space, and that all matter 
is divisible. But where does this divisibility end? We 
do not know. A single costal of a dark purple solid 
known as potassium permanganate will very perceptibly 
color several gallons of water. To clo this, it must be 
divided up until its particles are diffused throughout the 
entire volume of the water, or into so many parts that 
the numbers are beyond our comprehension. Though we 
cannot fix an absolute limit to the division of matter, we 
assume it to be the molecule. 

9. The Theory of Molecules. — A molecule is the smallest 
particle of matter that can exist alone, or in the case of a 
compound body, the smallest particle that can exist with- 
out destroying the identity of the substance. Thus, the 
smallest portion of common salt to be conceived of still 
contains the two elements mentioned, sodium and chlorine, 
and is a molecule. In the same way the smallest particle 
of water would contain both hydrogen and oxygen, and 
always in the same ratio. If, now, by any means we can 
break up these salt or water molecules into their two con- 
stituents, we no longer have a compound body, but two 
simpler substances or elements. 

10. Atoms. — These constituent parts of a molecule we 
call atoms. Even the molecules of the elements may con- 
sist of two or more atoms ; in fact, they usually do. 
Thus, it will be seen later that a molecule of oxygen con- 
sists of two atoms ; a molecule of chlorine of two atoms, 
and so on. It may be said, therefore, that all matter is 
divisible into molecules, and that these molecules are 



12 



MODERN CHEMISTRY 




composed of still smaller particles called atoms. From 
the above, it will be seen that the molecule of the com- 
pound body and its constituent atoms would be very dif- 
ferent, while the molecule of an elementary body and its 
atoms would be exactly alike in properties. This is illus- 
trated in Fig. 1. 

Here, a represents a molecule of water consisting of two atoms of 
hydrogen and one of oxygen. If by electricity we decompose this 
molecule, we shall no longer have water, but 
two elementary substances, hydrogen and oxy- 
gen. Likewise, c represents a molecule of 
common salt, and if this be decomposed, we 
shall no longer have salt, but two substances, 
sodium and chlorine. On the other hand, b 
represents a molecule of oxygen, and if this 
be decomposed, we shall still have oxygen, 
but in the atomic form, possessing the same 
physical characteristics as before. 

11. Physical Changes. — Matter exists in three condi- 
tions, solid, liquid, and gaseous, depending upon the rela- 
tion to each other of the intermolecular forces. When 
the cohesion existing between the molecules is consider- 
ably greater than the repellant forces which tend to drive 
them apart, the substance exists in the form of a solid. 
When the converse of this is true, and the molecules tend 
to be driven farther and farther from each other, the sub- 
stance is then in the form of a gas. When the repellant 
and attractive forces are about at an equilibrium, the 
substance exists as a liquid. It is with changes of molec- 
ular condition that physics deals, and the molecule is the 
basis of all physical phenomena. 

Experiment 1. — Hold in the flame of the Bunsen burner a piece 
of tin foil, an aluminum wire, or a narrow strip of zinc; or put any 
one of them into a clean iron spoon and heat. What change takes 






PHYSICAL CHANGES 13 

place in the physical condition of the metal? Let it cool again and 
state what happens. 

12. In this simple experiment two changes have taken 
place : one, the conversion of the solid into the liquid 
form; and second, the changing back again into the solid. 
We have changed the form of the substance and the 
arrangement of the molecules as regards one another, but 
the properties have remained the same. 

13. Another Molecular Change. — Likewise, when a cur- 
rent of electricity is passed through an electro-magnet, 
the armature for the time being becomes a magnet ; that 
is, its molecules have been rearranged, or so affected that 
they present the well-knowui phenomenon of attraction. 
When the circuit is broken, the armature loses its mag- 
netism ; in other words, its molecules have assumed their 
previous condition. When a body is heated, the molecules 
are set in more rapid vibration, and luminous bodies emit 
light because of this vibratory motion. 

14. So throughout the domain of physics w^e find that 
all phenomena concern the molecule and molecular changes. 
We are all familiar with the different conditions under 
which water exists, and we are not surprised at the state- 
ment that fog, clouds, rain, snow, and other forms are 
modifications of the same substance, the only difference 
being in the greater or less amount of stored-up heat 
energy. 

15. Substances exist in Three Forms. — At the same time 
we are not accustomed to think that nearly all substances 
may exist in the same three forms ; as liquids, solids, and 
gases. We are familiar with air in the gaseous condi- 
tion only ; yet if we reduce the temperature of it to about 
190° below zero Centigrade, it becomes a transparent 
liquid not very different in appearance from water, and 



14 MODERN CHEMISTBY 

at a still lower temperature it freezes or solidifies. In 
like manner carbon dioxide, an invisible gas thrown off 
from the lungs of all animals in breathing, if cooled, is 
first liquefied, and then by further reduction of tempera- 
ture is converted into a beautiful white crystalline solid, 
very closely resembling snow. 

16. Mercury and Carbon. — We are all familiar with 
mercury as a silvery white liquid, which may be boiled 
and converted into vapor at a moderately low temperature. 
On the other hand, in our most northern climates it fre- 
quently solidifies, and were the glacial epoch to return, 
under the rigors of that era we would know mercury, 
not as a liquid, but as a somewhat malleable solid closely 
resembling lead in its appearance and properties. To go 
still further, carbon, that we know best in the form of 
charcoal and coal, — one of the most refractory of sub- 
stances — in an electrical furnace may be fused and made 
to boil, apparently like water, while in the intense heat 
of the sun it is converted into vapor and exists in the 
atmosphere surrounding that body just as water vapor 
does in our own. 

17. All these are merely illustrations of changes in 
molecular condition, or are physical changes ; and the 
same substance in its different forms presents at all times 
essentially the same properties. 

Experiment 2. — Put into a small test-tube a crystal of iodine 
and warm gently. What becomes of the crystal? Does anything 
deposit farther up the tube? When the colored vapors have disap- 
peared, warm the tube higher up. What happens ? 

From the above experiment we see what has been stated 
before, that heat converts many substances into vapors ; 
and further, that when this heat is removed, they condense 
again into their previous condition. We learn, too, that 



and 



CHEMICAL CHANGES 15 

the physical condition in which bodies exist is not an essen- 
tial, but an incidental matter depending upon the ease or 
difficulty with which they are melted and vaporized. 

18. Chemical Changes. — Chemical changes, on the other 
hand, involve, not the molecule, but the constituent parts 
of the molecule, the atoms. In every chemical change the 
molecule is broken up and its identity destroyed, while 
the atoms formerly composing it recombine to form ne^\ 
and different substances. Hence, when any substance i^ 

nged chemically, or when two or more substances react 
n one another and undergo a chemical change, new 
bducts are formed, differing altogether in properties 
from the original substances. 

Experiment 3. — Put into a small test-tube a little mercuric oxide 
and heat, gently at first, then quite strongly. Continue this for sev- 
eral minutes. Thrust a pine splinter, having a spark upon the end, 
down into the tube. What happens? Is there any evidence of some- 
thing present different from air? What seems to be forming upon 
the cooler portions of the tube ? 

19. If we continue to heat long enough, the red oxide 
would entirely disappear, just as water does when boiled. 
In the latter case, however, the water might be condensed 
again as before, while in the present instance the mercuric 
oxide seems to be decomposed into two substances, — one, 
an invisible gas which caused the spark to burn brightly, 
the other a dark colored substance which condensed in 
small globules upon the cooler portion of the tube. Here 
we have a chemical change, one which caused the destruc- 
tion of the original substance, and produced therefrom 
two, differing very strikingly in their properties. 

20. Chemical changes are usually brought about by some 
physical agency, such as heat, electricity, light, percussion, 
etc. Innumerable instances of this kind of change might 



16 MODERN CHEMISTRY 

be given, but as such changes are to be studied throughout 
the science of chemistry, only a few will be noticed at 
present. As already seen, the substances produced are 
often very different from those used in obtaining them ; 
for example, two or more solids may unite in such a way 
as to form a gaseous body, or even a liquid ; two gases 
may form a solid or a liquid; while two liquids may 
combine to form a solid. Some of these will be illus- 
trated below. 



Experiment 4. — Rub vigorously together in a mortar a si 
crystal of potassium chlorate and an equal quantity of flowed 
sulphur. Into what are the two solids converted? How is i 
chemical union manifest? 



fcW^I 



21. A change similar to this is seen when gunpowder 
is exploded, causing the three substances, of which it is 
a mixture, to combine, with the formation of several 
gaseous products. 

Experiment 5. — With a few small crystals of potassium chlorate 

wrap in a piece of paper a bit of phosphorus ; fold the ends together 

carefully and strike with a heavy weight as shown in 

Fig. 2. What are the results? What is the nature of 

the products formed? 

CAUTION. — In preparing for this experi- 
ment, cut the phosphorus under water, dry 
quickly in the folds of a filter or blotting 
paper, and proceed as above. Small parti- 
cles of the phosphorus often fly to some dis- 
FlG * 2t tance, and a board should be set up to protect 

the clothing of the experimenter and of the members of 
the class. 

Experiment 6. — Mix thoroughly a small quantity of granulated 
sugar with an equal amount of potassium chlorate, well powdered, 
and put the mixture into an iron saucer. Now by means of a pipette 




CHEMICAL CHANGES 



17 



or glass tube drop upon it a little strong sulphuric acid. Notice that 
the sugar is burned, a part is converted into gaseous products, and a 
part left as a black residue. 

22. Chemical Change by Physical Agency. — In Experi- 
ments 3, 4, and 5, the different substances were caused to 
combine by friction, or percussion. That is, by physi- 
cal force, the molecules of the different substances were 
brought so close together that the chemism, or chemical 
Affinity, of the atoms in the unlike substances w^as greater 
tliK the cohesion among the molecules, and as a result 
^l^atoms rearranged themselves to form new compounds. 

This may also be illustrated by the following experiment : 

Experiment 7. — Put into a clean porcelain mortar a few crystals 
of potassium iodide, KI, and about the same amount of mercuric 
chloride, HgCl 2 . Rub them together for two or three minutes; notice 
that the two white compounds react with each other to form a bright 
red mixture. 





23. By the friction, the molecules of the mercuric chlo- 
ride and of potassium iodide are brought so close together, 
that the mercury and potassium atoms exchange places, 
forming mercuric iodide, which is bright red in color, and 
potassium chloride, which is a white compound. It is 
simply another striking illustration of the fact already 
stated that chemical action brings about a change in the 
atomic structure, and causes, therefore, the formation of 
substances very different from the original. 



18 MODERN CHEMISTRY 

24. Chemical Change by Chemical Agent. — In Experi- 
ment 6, the same results were brought about by the use 
of a chemical reagent, sulphuric acid. In all of these 
instances, except Experiment 7, the solid substances used 
were largely converted into gaseous products, and in all 
entirely different from the original. To show that two 
gases may combine to form a solid, perform the following 
experiment : — 

Experiment 8. — Pat into a beaker a few drops of hydroc 
acid, cover with a sheet of cardboard or glass, arid allow it to st; 
few minutes. Into another beaker put a few drops of strong ai 
nium hydroxide, and cover in the same way. Presently, inver 
first beaker over the second, then remove the cards, so that the two 
gases which have filled the beakers may come into contact with each 
other. Notice the heavy white fumes that form. These are a white 
solid compound, known as sal ammoniac. 

Experiment 9. — To show that two liquids by reacting upon each 
other may produce a solid. Numerous instances of this will be seen 
from time to time. Powder a gram or two of alum and an equal 
amount of ferrous sulphate, put both into a test-tube and add about 
10 cc. of water or just enough to dissolve them readily. You will now 
have a nearly colorless solution. To this add slowly a little strong 
ammonium hydroxide. You will obtain a greenish colored, jelly-like 
solid, with scarcely enough water present to enable the precipitates to 
be poured from the tube. 

25. Mixtures. — We have had the term compound body 
defined, and we must be careful to distinguish between a 
compound and a mixture. In the latter, the two or more 
substances used are not necessarily in any definite propor- 
tion, nor are they united with each other. Furthermore, 
as a rule, mixtures may be very easily separated by purely 
mechanical means. 

Thus, we may put together sand and common salt in 
any proportion whatsoever; they do not react with each 
other to form a new compound, but are still sand and salt 



hioric 
rt~cnw 



I 



MIXTURES 19 

just as before being mixed. They may also be very easily 
separated and the salt recovered. Let the student suggest 
a method for doing this. 

26. Gunpowder. — Gunpowder is a very familiar mix- 
ture, consisting of three substances which may be easily 
separated. Upon exploding, however, these three combine 
to form several compounds from which the original could 
be recovered only by difficult and expensive methods. 

Experiment 10. — To separate gunpowder into its constituents. 
Put a gram or two of gunpowder into an evaporating dish and add a 
few cubic centimeters of water. After warming gently a few minutes, 
filter out as directed on page 00. Transfer the clear filtrate to an 
evaporating dish and boil slowly to dryness. While you are waiting 
for this, transfer the black residue upon the filter paper to a beaker 
and add a little carbon disulphide. Shake for a moment or two, and 
then decant or filter off this clear solution. Without heating, let it 
evaporate to dryness in a watch crystal or evaporating dish. What 
familiar substance is left? What was the appearance of the solid 
obtained by boiling down the solution in water? What colored resi- 
due was left after treating with carbon disulphide? Can you name 
the three substances? 

Experiment 11. — To show the difference between a mixture and 
a compound of the same two elements. Put together in any proportion 
a small quantity of sulphur and of iron filings. Mix them thoroughly. 
What is the resulting color? See whether you can remove the filings 
by means of a magnet ; owing to the presence of some moisture a 
little sulphur may adhere to the filings. State your results. Put the 
filings back into the sulphur and again mix them well. Now add 
1 or 2 cc. of carbon disulphide, shake for a moment or two, and 
decant thoroughly upon a watch crystal or into an evaporating dish. 
Let the clear solution dry without heating, as it is very inflammable. 
What do you obtain ? Have you effected a separation of the two ? 

Again mix in about equal proportions by volume, or in the ratio 
of 32 to 56 by weight, sulphur and iron filings ; put them into the 
smallest test-tube you have and heat, first gently and then very 
strongly, until a bright red glow seems to go through the entire mass. 
After thus heating for two or three minutes, allow the tube to cool, 



20 MODERN CHEMISTRY 

and remove the contents ; if necessary, break the tube. How does 
the color differ from that of the mixture at the beginning? Powder 
the mass, and attempt to separate the filings from the sulphur by 
means of a magnet as before. Can you do this? Treat a part of 
the dark powdered mass with carbon disulphide and determine 
whether you can thus remove the sulphur as you did from the mix- 
ture. State your results. From this experiment we may note several 
differences between the mixture of iron and sulphur that we had 
before heating and the compound afterward. What are they? 



SUMMARY OF CHAPTER 

Introduction. 

Some theories of matter — Old — Illustration. 

Present theory. 

Definition of term element. 
Compound bodies. 

Definition. 

Illustration. 
Divisibility of matter. Chemical changes. 

Difference between molecule How different from physical, 

and atom. Several experiments to illustrate. 

Illustration. Mixtures. 

Physical changes. How different from elements and 

Experiments to illustrate. compounds. 

Name several others. Examples of mixtures. 

What is a physical change ? Method of separating. 



CHAPTER II 

VALENCE 

1. What is Valence ? * — If we notice some of the com- 
mon compounds of hydrogen, which we shall study, we 
shall see that different elements unite with a different 
number of hydrogen atoms. Thus, chlorine combines with 
one atom of hydrogen, oxygen with two, nitrogen with 
three, and carbon with four, as shown in the following 
compounds : — 

Hydrochloric Acid .... HC1 

Water H 2 

Ammonia H 3 N 

Marsh Gas ...... H 4 C 

2. The four elements, chlorine, oxygen, nitrogen, and 
carbon, have the power of combining with one, two, three, 
and four atoms of hydrogen, respectively, and we speak 
of them as having a valence of one, two, three, and four. 

I By valence or quantivalence we 
mean the power any element has ^ , , t 

of holding in combination the N< • » ■ 

atoms of another element taken 

as a standard. / This standard, primarily, as shown above, 



* If in the judgment of the teacher the subject of valence can be more 
easily grasped by the student later in the course, this chapter may be 
deferred until after the study of carbon and its compounds. 

21 



22 MODERN CHEMISTRY 

is hydrogen, and by it the valence of other elements is 
measured or determined. It may be illustrated in this 
way : suppose the first line represents the combining 
power of hydrogen, which is our standard. Then with 
this " yard stick " we will measure the combining power 
of the other elements. In water, H 2 0, the valence of 
the oxygen atom is determined by applying the "yard 
stick," and is seen to be two ; in NH 3 the standard is used 
three times, and the valence of the nitrogen atom is three. 
In the same manner the valence of the carbon atom is 
determined as four. 

3. Suppose, however, hydrogen did not combine with 
carbon, could we still determine its valence? We are 
familiar with the compound, carbon dioxide. In this 
molecule, C0 2 , the oxygen atom is used twice with the 
carbon atom, hence the latter must have twice the combin- 
ing power of the oxygen. This has already been shown 
to be two, hence carbon would be four. 

To illustrate roughly, we sometimes speak of the atoms 
as having a certain number of " bonds " or " poles of 
attraction," represented as below : — 

(sy (c^p @=» Atoms with one " bond." 

^ Atoms with two " bonds." 

*@R C (^) D Atoms with three " bonds." 

c(c\> ^(sih Atoms with four " bonds." 

From this illustration it will be seen that an atom in 
the second group in combining would have two bonds by 



VALENCE 



23 



which to hold the two bonds of two individual atoms of 
the first. 






Lead Chloride. 



Mercuric Bromide. 



Water. 



(£}<*><£) <g><^Hg) N and Sb both have valence 

of three. 



Ammonia. 



Antimony Chloride. 



(5Xc><£) (£>sgg© 



C and Si, valence of four. 



Marsh Gas. 



Silicon Dioxide. 



4. Such atoms as combine with one of hydrogen or its 
equivalent are said to be univalent, or are sometimes called 
monads; those which combine with two of hydrogen are 
bivalent, or dyads ; with three, trivalent, or triads ; with 
four, quadrivalent, or tetrads; with- five, quinquivalent, or 
pentads. 

5. Variation in Valence. — In studying a number of the 
compounds of any element it will be noticed that while 
the valence of the element in most of them is the same, 
there will be some compounds which show it to be dif- 
ferent. Many of these are believed to be merely apparent^ 
exceptions and may be readily explained ; while others, 
as yet not thoroughly understood, may be real variations. 
For example, the oxygen atom is always regarded as biva- 
lent, yet we shall meet with the compound, hydrogen 
peroxide, H 2 2 , in which oxygen is apparently univalent. 



24 MODERN CHEMISTRY 

It is believed, however, that the atoms have an arrange- 
ment in the molecule which may be represented thijs : — 

^ or ®X£><£><® 

This simply means that one bond of each atom of 
oxygen is held by a bond of the other. In a similar way 
the apparent double valence of a great many other ele- 
ments is explained. Thus copper and mercury, ordinarily 
bivalent, also form the compounds Cu 2 Cl 2 , and Hg 2 Cl 2 . 
But these are exactly parallel to the case above. 



?® ^"® /WS> ^>"® 



Cul 

\ci) (^<a) TQ) @H«) 

Bivalent. Apparently Bivalent. Apparently 

Univalent. Univalent. 

6. Again we shall study the two compounds of carbon, 
CO and C0 2 , the first of which would indicate a valence 
of two for the atom, while in the second it would be four. 
The second is believed to show the true valence, and 
carbon monoxide is regarded as an unsaturated compound, 
that is, one in which the valence of the atom is not satisfied, 
or one in which a part of the bonds is not held by any 
other element. We may represent it thus : — 



Saturated Unsaturated 

Compound. Compound. 

This theory is accepted because carbon monoxide very 
readily takes up one more atom of oxygen and forms the 
dioxide. 



VALENCE 25 

7. Double Valence. — In the examples of double valence, 
noticed aboVe, the irregularity is only apparent. There are 
many cases, however, in which all the indications thus far 
would show that the valence of the atom is variable. 
Thus we have said the nitrogen atom is trivalent, and 
this is so in ammonia and nitrogen trioxide ; but we shall 
also meet with nitrogen pentoxide, N 2 5 , in which the 
valence is five ; the monoxide, N 2 0, wherein it is appar- 
ently one, etc. There are many such variations that will 
trouble the student, but for our present work we shall 
need, as a rule, to aid us in writing formulae and reactions, 
only a knowledge of the ordinary valence of the atoms. 

8. Valence of Groups or Radicals. — We shall find also 
that many groups of atoms react in the same way as 
individual atoms ; such groups are called radicals. They 
have combining power or valence just as individual atoms 
have. Thus when sulphuric acid reacts with zinc, we shall 
find that the group (S0 4 ) is not broken up ; the same is 
true in hundreds of other instances. As it is combined 
with two atoms of hydrogen in H 2 S0 4 , and always does 
combine in the same way, we say its valence is two ; this 
may be shown graphically in this way. 

°) ®><&J$) 

(S0 4 ) Group, showing Sulphuric Acid, showing all honds 

two bonds unused. of the (S0 4 ) group saturated. 

While we cannot prove that such is the arrangement of 
the atoms in the group, still it is believed to be true ; at 
any rate it serves to illustrate how the valence of the 
group is two. In the same way we would show the 
valence of any other radical. In sal ammoniac, NH 4 C1, we 



26 MODERN CHEMISTRY 

find the group (NH 4 ) in combination with one atom of 
chlorine, hence its valence is one. Then if the radical 
(NH 4 ) combines with (S0 4 ), it must be used in the 
proportion of two of the former to one of the latter, thus 
(NH 4 ) 2 S0 4 . 



sY mp<a 



Sulphuric Acid. Ammonium Ammonium 

Chloride. Sulphate. 

Or we may show the same facts in this way : — 

(S) 






W% ■ ®f® 



Ammonium, showing Ammonium Sulphate, showing how (S0 4 ) 

one bond unused. can unite with two of (NH 4 ) . 

Exercise in Valence.* — Applying the principles set forth in the pre- 
ceding paragraphs, let the student write the formulae for the fol- 
lowing : — 

when Ba unites with I, CI, Br, O, N0 3 , Cr0 4 . 

when Na unites with O, S, CI, C10 3 , Si0 4 , S0 4 . 

when Cu unites with CI, S, S0 4 , HO, O, I. 

when NH 4 unites with I, P0 4 , S0 4 , S, HO, Br. 

when Bi unites with CI, O, S, N0 3 , S0 4 . 

* Let the teacher give further exercises until the student can write with 
assurance the compound resulting from the union of any two of these 
elements or radicals, 



below : - 





VALENCE 




mce of 


each of the 


above and 


certain othe 


Monads 


Dyads 


Tkiads 


Tetkads 


I 


Ba 


Sb 


c 


Br. 


Zn 


Bi 


Si 


CI 


Ca 


As 


(Si0 4 ) 


F 


Sr 


(P0 4 ) 




Na 









K 


Hg 






Li 


Cu 






(NH 4 ) 


Fe 






(N0 8 ) 


S 






(C10 3 ) 


(S0 4 ) 






(HO) 


(Cr0 4 ) 







27 



9. Equations. — We shall find later on that we are 
required often to write equations which represent the 
chemical action that takes place when two or more com- 
pounds are put together. We shall learn that chlorine 
and hydrogen, if put together in equal proportions, by a 
spark may be made to combine and form a compound 
body; we would represent this briefly as follows : — 




or, 



H 9 + CL = 2 HC1. 



So if we put two compounds together, we would express 
the chemical action by a similar equation, thus: — 





or, 



HC1 + AgN0 3 = HN0 3 + AgCl. 



28 MODERN CHEMISTRY 

It will be noticed that the first elements in the compound 
bodies have simply interchanged places. So, if we bring 
together ammonium hydroxide, NH 4 OH, and sulphuric 
acid, H 2 S0 4 , we shall find a similar exchange,, thus : — 

2NH 4 OH + H 2 S0 4 = (NH 4 ) 2 S0 4 + 2 H 2 0, 

in which we are compelled to use two parts of the NH 4 OH 
because the group (NH 4 ) has a valence of only one, and 
there are two atoms of hydrogen in the acid to be dis- 
placed. 

£)„ 

{si 



Exercises.* — In the same way let the student show by simple 
equations the changes which take place when the following com- 
pounds are put together : — 

NaBr + AgN0 3 = HgCl + KI = 

Zn -f H 2 S0 4 = NH 4 OH + HN0 3 = 

NH 4 OH + HC1 = . Pb(N0 3 ) 2 + 2 HC1 = 

BaCl 2 + H 2 S0 4 F 

SUMMARY OF CHAPTER 

Valence — Meaning of term. 

Illustrations. 
Classification of elements as to valence. 

Synonymous terms for univalent, etc. 
Variation in valence. 

Apparent — Illustrations of. 

* Let the teacher add as many others as may seem necessary. 



WATER 29 



Unsaturated compounds. 

Real — Illustrations of. 
Valence of radicals. 

Illustrations. 
Application of a knowledge of valence. 

Writing formulae of compounds. 

Writing equations. 



CHAPTER III 

WATER : H 2 

1. Its Abundance. — Water is one of the most abundant 
substances known ; it covers about three-fourths of the 
surface of the earth, besides existing in vast quantities in 
other forms. In the arctic regions in the form of an ocean 
of compressed snow it covers the entire surface of the land 
to a depth of many feet; in a similar form it caps all the 
loftiest mountain peaks from which great rivers of ice 
flow down the valleys until they are melted at the snow- 
line. In the form of vapor it exists in the atmosphere, 
invisible except when condensed in fogs, clouds, etc. In 
any given locality this moisture in the air varies largely 
at different times, but not often is there more than sixty- 
five per cent of what the air is able to hold. Even with 
this amount, however, it has been estimated that were the 
vapor in the air all condensed, it would form over the 
surface of the entire earth a layer of water five inches 
deep. 

2. The human body rs about sixty per cent water, and 
daily throws off through the pores and from the lungs 
over three pounds of moisture. Many vegetable articles 
of food contain eighty to ninety per cent of water, and 
some even more. 



30 MODERN CHEMISTBY 

3. Water of Crystallization. — Water also exists in an- 
other form not so familiar as those already mentioned; 
that is, tvater of crystallization. A great many compounds 
in solidifying from their aqueous solutions take up a con- 
siderable amount of water. This does not exist in a free 
state like water in the pores of a sponge, or in a piece of 
soft wood that has been submerged for some time, but is 
in combination — crystallized in with the molecules them- 
selves. Such substances in crystallizing cause the disap- 
pearance of a considerable amount of water, which may, 
however, usually be obtained again by subjecting the body 
to a greater or less degree of heat. Some astronomers 
even believe that the absence of water upon the moon may 
be accounted for by the fact that such bodies as those 
mentioned have taken it all up in crystallizing from their 
aqueous solutions. An idea of the amount contained by 
such substances may be gained from the following experi- 
ments. 

Experiment 12. — Put into a test-tube a crystal of native gypsum 
and heat in the Buusen flame. What do you see depositing upon the 
cooler portions of the tube ? How is the crystal of gypsum affected ? 
Repeat the experiment, using borax or alum instead of gypsum, and 
state results. 

Experiment 13. — Expose to the air for several hours a crystal of 
Hn. ferrous sulphate. Notice its appearance before the exposure ; how- 
has it changed in the air ? 

4. Efflorescent Substances. — Many such substances as 
ferrous sulphate and copper sulphate, upon being exposed 
to an atmosphere more or less dry, give up all or part of 
their water of crystallization ; at the same time they 
usually change in color a,nd crumble to a powder. The 
process is the same as when the substances are heated, but 
not so rapid. By adding water to them the color is 



\ 



WATER 31 

usually restored, and they crystallize as before. Such 
substances as these that give up their water of crystalliza- 
tion to the air are said to be efflorescent. 

5. Deliquescent Substances. — There is another class of 
substances which have the power of abstracting moisture 
from the air or surrounding bodies, and of dissolving them- 
selves either in whole or in part in this moisture. Such 
are called deliquescent bodies. A familiar example of 
these is common lime, which on being expossd to the air 
gradually takes up moisture, crumbles to a powder, and 
becomes " air-slaked." Another noted example is phos- 
phorus pentoxide, a white solid formed when phosphorus 
is burned in the air or oxygen ; also calcium chloride and 
caustic potash. 

Experiment 14. — Put into a dry evaporating dish or beaker a 
small lump of fused calcium chloride and allow it to stand several 
hours or over night. Xotice how it has changed. In the same way 
expose a small piece of caustic potash. Notice how rapidly it changes. 
Only a few minutes will be necessary in this case. 

Common salt, stick candy, and some forms of taffy are 
very familiar examples of deliquescent bodies. 

6. Distinguishing Characteristics of Water. — In the pure 
state, water is practically colorless, but when of great 
depth it is seen to be of a blue color. It is odorless and 
tasteless, but we are so accustomed to drinking impure 
water that when we use that which is distilled, or 
perfectly pure, it tastes " flat," just as unseasoned food 
does to those who are habituated to the use of salt, 
pepper, and other condiments. Pure water, on being 
evaporated to dryness, leaves no residue whatever, and 
this, in connection with the fact that it affects vege- 
table coloring matter in no way, is one method of test- 
ing it. 



32 MODERN CHEMIST RY 

7. Solvent Powers of Water. — Pure water is seldom 
found, owing to its great solvent powers. To a greater 
or less extent it may be said to be almost a universal 
solvent. Even glass and similar substances immersed in 
water show appreciable loss after a considerable length of 
time. From this property result the various kinds of 
" hard " or mineral waters, medicinal, saline, etc. It is 
owing to the solvent powers of water, and the fact that 
evaporation leaves all mineral matter behind, that the 
ocean contains such vast quantities of different kinds of 
salt. 

8. For example, in a hundred pounds of sea water, there 
are over three pounds of solid matter ; of this the greater 
portion is common salt, but compounds of magnesium 
and calcium in the form of what are usually known as 
epsom salts and gypsum also occur. It has been estimated 
that if the ocean were of an average depth of one thousand 
feet, the common salt in solution would occupy a space of 
about three and a half million cubic miles, or a volume 
more than five times as great as that of the Alps. On 
this basis, if the depth of the ocean averages what is now 
claimed for it, the amount of salt surpasses in bulk our 
greatest mountain ranges. 

9. Composition of Water. — By many of the ancients, 
water, along with fire and air, was regarded as an element; 
but about 1800 a.d. it was proved to be a compound body. 
There are two methods of proof, which taken together 
are quite conclusive. 

The first proof is by Electrolysis. 

Experiment 15. — Fill the tubes of the electrolytic apparatus 
shown in Fig. 3 with water slightly acidulated with sulphuric acid, 
the latter being added simply to increase the conductivity. Then 
connect the platinum electrodes with a strong battery. As the 



WATER 



33 




current passes through the water, bubbles of gas will be seen rising 
from the two strips of platinum, P, and 
from one of them considerably faster 
than from the other. It will be found 
that twice as much gas collects in one 
tube as in the other. These two gases, 
we shall learn before long, are hydro- 
gen and oxygen. Open the stop-cock 
of the tube containing the greater 
amount of gas, and hold a lighted 
match to it ; notice that it burns with 
a very pale flame. Test the gas in the 
other tube, using a pine splinter with 
a spark upon the end ; notice that it 
bursts into a flame. Fig. 3. 

10. The second proof is by Synthesis. 

Experiment 16. — Put into the eudiometer, Fig. 4, 8 cc. of oxy- 
gen, and twice as much or more hydrogen, the instrument being 

already partly filled with mercury, M. 
Hold the thumb over the open end and 
pass a spark from a galvanic battery. 
The two gases will combine with ex- 
plosive force, producing water in the 
form of vapor. If the proportions were 
exactly two of hydrogen to one of oxy- 
gen, when the apparatus has become 
cool, there will be no gaseous residue, 
showing that the two unite in this pro- 
portion to form water. A in the figure 
is a cushion of air left to break the force of the explosion. 




Fig. 4. 



11. Conclusions. — From the above experiments we 
learn that water is the result of the union of two invisible 
gases, one of which burns with a pale flame, the other of 
which causes various substances to burn vigorously. We 
see also that from the union of the two gases, which to- 
gether form a very explosive mixture, there results an 



34 



MODERN CHEMISTRY 



exceedingly stable compound, which not only does not 
burn, but which has the power of quenching thirst and of 
overcoming the greatest fires. These two gases were 
given the names, hydrogen and oxygen. 

12. We noticed also in the proof by analysis, that 
the hydrogen was given off in volume double that of the 
oxygen, and further, that in mixing the two gases for the 
synthetic proof we caused them to unite in the same ratio. 
From theso experiments we may conclude that the com- 
position of water by volume is two parts of hydrogen to 
one of oxygen, a fact which we represent by the expression 
H 2 0. 

13. Analysis by Other Methods. — The analysis of water 
may be effected by means other than electricity. For 
example, if a current of steam is made to pass through a 
tube containing charcoal or coke heated red hot, the steam 
is decomposed ; the oxygen combines with the carbon of 
the charcoal, forming an oxide of carbon, and at the same 
time the hydrogen is set free. 

14. Synthesis by Other Methods. — In a similar way 
the synthesis of water may be effected. If a current of 
hydrogen is passed through a tube containing some me- 
tallic oxide, heated to redness, for example, copper oxide, 
the oxygen is removed from the compound by means of 
the hydrogen, and water is formed and may be collected. 

P^xperiment 17. — Into a of the small bulb-tube put a little black 

oxide of copper, and weigh 
both tube and oxide care- 
fully. Next fill a U-shaped 
tube with lumps of calcium 
chloride, weigh and quickly 
connect with the other tube. 
Now pass a current of hy- 
Fig. 5. drogen, generated as on 




HjzSOj. 



WATER 35 

page 39, over the copper oxide, heated to redness. The hydrogen 
should first be dried by passing through sulphuric acid or over calcium 
chloride. After some time, disconnect the apparatus, and weigh the 
U-tube ; the gain in weight will represent the amount of water pro- 
duced. When the bulb-tube is cool, weigh it : the loss will represent 
the amount of oxygen removed. Subtracting the weight of the oxygen 
from the weight of water found will give the amount of hydrogen. 
Allowing for errors, this should give eight parts of oxygen to one of 
hydrogen, by weight. 

From this experiment we are able to conclude as to the quantita- 
tive composition of water, just as by the others we learned of the 
volumetric. The action of hydrogen in thus removing oxygen from 
an oxide is called reduction. 

Water | By volume : H y dr °g en > 2 ; Oxygen, 1. 
( By weight : Hydrogen, 1 ; Oxygen, 8. 



SUMMARY OF CHAPTER 

Water — Various forms in which it occurs. 
Water of crystallization. 
Meaning of term. 

Examples of substances containing it. 
Proof of its presence by experiment. 
Efflorescent substances. 
Deliquescence — Meaning of term. 

Illustrations. 
Some special characteristics of water. 
Composition of water — Proof of. 

a. By analysis — Details of work. 

Apparatus used. 

b. By synthesis — Explanation of process. 

Drawing of apparatus. 
Composition by weight. 
Proof by experiments. 



CHAPTER IV 

HYDROGEN : H = 1 

1. History. — The term hydrogen is from two Greek 
words, which mean water producer, and the gas is so 
named because this element enters so largely into the 
composition of water. It was first isolated in quantities 
sufficient for experiment by Cavendish in 1766, and on 
account of its combustibility was called by him inflam- 
mable air. 

2. Where found. — Hydrogen is seldom found uncom- 
bined, though its chemical affinity for most substances is 
not very marked. It exists abundantly in composition — 
water being the most important example ; it enters into 
nearly all organic compounds ; it is given off, together 
with other gases, by some volcanoes ; and by the spectro- 
scope we know that it exists in the atmospheres of the sun 
and of some of the stars. 

3. Methods of obtaining Hydrogen. — We have seen 
already in Experiment 15 that hydrogen may be ob- 
tained by the electrolysis of water. This gives a very- 
pure gas, but does not produce it rapidly enough for 
ordinary experimental purposes. Just as electricity has 
the power of decomposing water, so do certain metals. 
When iron is exposed to moisture, we say it rusts; in 
reality, it takes up a certain amount of oxygen from the 
water and sets free a corresponding amount of hydrogen.* 

* Rust is an oxide of iron ; that is, a compound of iron and oxy- 
gen, represented by the formula Fe 2 3 . The chemical change which 

36 



HYDROGEX 



37 



4. Decomposition of Water. — Again, there are some 
metals, like calcium, a constituent of common limestone, 
which have the power of decomposing water at the boiling 
point, setting free a part of the hydrogen and forming at 
the same time a compound, such as lime water. The 
chemical action may be expressed by the following 
equation : — 

Ca + 2 H 2 = Ca (OH) 2 + H a . 

5. There are two common metals, sodium and potas- 
sium, which decompose water rapidly at ordinary tempera- 
tures. Of these, the second acts much more violently, 
generating almost instantly sufficient heat to ignite the 
hydrogen given off from the water and volatilizing a por- 
tion of the metal itself. This is seen in the violet color 
which is imparted to the flame. That sodium is setting 
free a combustible gas in the same 
way may be shown by bringing a 
lighted match close to the metal, 
when the hydrogen will be ignited as 
it was spontaneously with potassium. 

Experiment 18. — Fill a test-tube with 
water and invert it over a trough or basin of 
water, as show T n in Fig. 6. Pat into a wire 
gauze spoon, or wrap in a piece of flexible wire Fig. 6. 




probably takes place when iron is thus exposed may be expressed as 

2 Fe -f 6 H 2 = 6 H + Fe 2 3 , 3 H 2 0, 

in which Fe 2 03 is rust. Likewise, if iron filings be heated red hot, and 
a current of steam slowly passed over them, the filings take up the oxygen 
from the steam and are converted into an oxide of iron, Fe 3 4 , differing 
somewhat from rust, while hydrogen is set free. The following equation 
expresses the chemical changes that take place : — 

3 Fe + 4 H 2 = Fe 3 4 + 4 H 2 . 



38 



MODERN CHEMISTRY 



cloth, a small piece of sodium and hold under the mouth of the tube. 
Bubbles of gas will rapidly form, will rise into the tube and displace 
the water. Test the gas obtained to see whether it acts as did the 
hydrogen obtained by electrolysis in Experiment 15. Does it seem to 
be the same kind of gas? Sometimes, before putting the sodium 
into water, it is treated with a small quantity of mercury, whereby 
the rapidity of the action is greatly decreased. 

6. Caustic Soda. — The chemical change which has 
taken place in the above experiment may be expressed 
as follows : — 




or, as it is usually and most simply written, — 

H 2 + Na = NaOH + H. 

The graphic equation above shows that in each molecule 
of water one atom of the hydrogen is replaced by one of 
sodium, represented by Na, and that thus a new compound, 
NaOH, called caustic soda or sodium hydroxide, is formed. 
In other words, the water molecule becomes one of caustic 
soda, thus : — 



becomes 



7. Tests to show Change. — We can easily prove that 
the solution has acquired new and very different proper- 
ties. Dip into it a strip of red litmus paper, and the 
paper is instantly turned blue ; or add to it one drop of 





HYDROGEN 



39 



a solution of phenol phthalein, and the whole will turn a 
beautiful red. 

8. Other Methods of obtaining Hydrogen. — The above 
methods of obtaining hydrogen, while of interest, are too 
expensive where considerable quantities are needed for 
experimental work. All acids contain hydrogen, and just 
as some metals decompose water, so certain others act with 
acids. . 

9. Laboratory Method. — In obtaining hydrogen for 
laboratory purposes this is the plan usually pursued. The 
metal used is generally iron or zinc, and the acid, sulphuric 
or hydrochloric. 

Experiment 19. — Fit to a flask a cork doubly perforated. Through 
one of the holes insert a delivery tube, and through the other a this- 
tle tube which extends 
nearly to the bottom of 
the flask. Put into the 
flask several pieces of 
granulated zinc made by 
pouring the metal in a 
molten condition into 
cold water. Add water 
until the zinc is nearly 
covered, and then pour 
in slowly a small quan- 
tity of strong sulphuric 
acid. After allowing the 
first gas which comes 
over to escape, because it is mixed with air, collect several bottles over 
water as described on page 362 and preserve for experiments a little 
later. The action may be hastened by adding a little copper sulphate 
to the flask a few minutes before the acid is introduced. 




Fig. 7. — Hydrogen Apparatus. 

p = pan. g = generating flask. 

t = thistle tube. r = receiving flask. 

d = delivery tube. 



10. Results of the Experiment. — The chemical change 
that takes place in the above preparation of hydrogen may 
be graphically represented as follows : — 



40 



MODERN CHEMISTRY 




From this it is seen that the atom of zinc has driven out 
or replaced two atoms of hydrogen in the sulphuric acid, 
H 2 S0 4 , and has entered into combination, while the two 
hydrogen atoms have been set free and now exist as a 
hydrogen molecule ; further, it is seen that the sulphuric 
acid has been converted into a compound known as zinc 
sulphate. Usually we express these facts thus : — 

Zn + H 9 S0 4 == ZnSO, + H„. 



Zinc sulphate is a white compound which collects upon 
the zinc, and would soon cover it so completely as to stop 
the chemical action ; the water, however, being added dis- 
solves it as fast as formed, and leaves a clean surface 
exposed to the acid. 

11. Method of obtaining Large Quantities. — When hydro- 
gen is desired in very large quantities, as in filling balloons, 
iron, being cheaper, is used instead of zinc. The gas thus 
obtained is somewhat less pure, but it is not on this 
account specially objectionable. Large vessels or retorts 
are used, which are lined with lead, a metal which is not 
affected by dilute sulphuric acid. The hydrogen obtained 
is passed through water and lime to purify it, after which 
it is transferred to the balloon. 

12. Mond's Method. — This method, although it has 
thus far been used only to a limited extent, promises to 
give satisfaction. We have seen that when steam is 



HYDROGEN 



41 



passed over red-hot charcoal the former is decomposed 
just as when passed over red-hot iron (page 37) and two 
similar products are formed, both gases ; thus, — 

H 2 + C = H 2 + CO. 

The apparatus may be represented conventionally as 
follows : — 




Fig. 8. — Mond's Method. 



F is a furnace, B the boiler in which the steam is generated, C a 
tube containing lumps of coke which are heated red hot by a gas 
furnace beneath, Ni a tube containing powdered metallic nickel, L an 
apartment containing lime water. The steam passing through C is 
decomposed as stated above; the mixture of hydrogen and carbon 
monoxide formed here passes over the nickel, also heated red hot. In 
this tube a part of the carbon unites with the nickel, and carbon 
dioxide is formed. This mixed with the hydrogen passes on through 
the " washer " L, containing lime water, which absorbs the carbon 
dioxide, leaving the hydrogen comparatively pure. The reactions in 
the different parts of the process may be shown as follows : — 

H 2 + C = H 2 + CO (in the coke tube). 
H 2 + 2 CO + Ni = NiC + H 2 + C0 2 (in the nickel tube). 
H 2 + C0 2 + Ca(OH) 2 = H 2 + CaC0 3 + H 2 (in the " washer "). 

The nickel carbide formed is readily converted back again into 
metallic nickel by heating in the air, so that it may be used over and 
over. 



42 



MODERN CHEMISTRY 




13. Characteristics of Hydrogen. — The following experi- 
ments will illustrate well the most striking peculiarities 
of hydrogen : — 

Experiment 20. — Remove one of the bottles of hydrogen from 

the water, keeping it inverted, and thrust up into it a burning candle. 

Notice whether the candle continues to burn in the gas ; notice also 

what happens as you remove it again. Can you see anything burning 

at the mouth of the bottle ? 

Experiment 21. — To show the lightness of hydrogen. Bring a 

bottle of the gas, a, mouth downward, up 

close to another inverted bottle, b, of about 

the same size. Then gradually tip the 

hydrogen bottle, a, as shown in Fig. 9, just 

as you would pour water from one vessel 

into another, only in a reverse order. 

Now test both bottles to learn which con- 
Fig. 9. — Upward Decantation. , . ., , , D , , ,, ,, 

tarns the hydrogen. State the results. 

Experiment 22. — Start the generator again, and replace the de- 
livery tube with one which has been drawn to a fine jet. Let the 
gas flow a few minutes until the air is all expelled, and then ignite it. 

14. CAUTION. — A mixture of air and hydrogen? is very 
explosive, and before lighting the jet a towel should be 
wrapped about the generating flask. It will do equally 
well to inclose the flask in a pasteboard box as shown in 
the figure. 

When first lighted, how does the hydrogen burn? How does it 
soon change? This is due to the sodium in the glass, which colors 
the flame. A burning jet of hydrogen is some- 
times called the "philosopher's lamp." Does the 
gas burn with much heat? Hold a clean dry 
bottle or test-tube over the, flame. Do you see 
any deposit forming upon the upper part of the 
tube? What is it? Now try several tubes and bot- 
tles of different sizes in the same way ; notice the 
different pitch of the " singing tones " produced. 
Fig 10 — Hvdiwen When the tube is thus sounding, notice the flame 
Jet in Box. carefully. Can you explain the tones produced ? 




HYDROGEX 43 

Experiment 23. — Allow a jet of hydrogen from a generator work- 
ing rapidly to strike against a platinum sponge. State the results. 

Experiment 24. — The hydrogen pistol shows the explosiveness 
of a mixture of air and hydrogen.* Load the pistol by pouring into 
it a small bottle of hydrogen, as shown in Experiment 21, and fire by 
bringing a flame to the touch-hole. A loud explosion should follow. 

Experiment 25. — Hydrogen soap-bubbles — to show lightness of 
hydrogen. For success in this experiment a good soap solution is 
necessary. To a little soft water add a few shavings of castile or 
other good soap, and when dissolved add about one-third as much 
glycerine as soap solution. Shake well. Xow attach to a delivery 
tube, from which is flowing a current of hydrogen, a clay pipe, or even 
an ordinary spool ; dip into the soap solution, and let the bubble form 
in the usual way. Detach from the pipe by a gentle jerk and notice 
whether the bubble rises or falls. Touch a light to one of the bubbles. 
What happens? This experiment is sometimes made more striking 
by filling the bubble with a mixture of hydrogen and oxygen, which, 
when touched with a flame, explodes violently. 

15. CAUTION. — The greatest care must be taken to avoid 
bringing the flame near the deliver}' tube, lest the whole 
mixture be exploded with serious results. 

Experiment 26. — To show the presence of hydrogen in oils, alco- 
hol, etc. We have already seen that when hydrogen burns, water is 
produced. This is true whether we have hydrogen free, or in the 
form of compounds. Light a small candle and hold over it a cold 
beaker. Xotice the water condensing upon the cooler portions of the 
beaker. In the same way try a small spirit lamp. State the results. 

16. Conclusions from our Work with Hydrogen. — By 
the above work with hydrogen we have learned that it is 
a colorless gas ; is without odor if pure, and very light. 

* The pistol may easily be made from a small tin can. With an awl 
punch a hole in the side of the can near the bottom, and for a bullet use 
a cork snugly fitted to the mouth of the can. When a light is brought to 
the touch-hole, several seconds may elapse before the explosion follows, 
but the experiment almost invariably succeeds. . 



44 



MODERN CHEMISTRY 



Its density is but little more than one-fifteenth that of 
air. It is this which causes it to diffuse so rapidly, and 
renders it valuable for filling balloons. A liter of the gas 
weighs .0896 g. It is very inflammable, and burns with 
a pale, almost non-luminous flame. As noticed above, the 
hydrogen flame from a glass jet has a yellow color, but this 
is due to a compound of sodium in the glass, just as the 
hydrogen arising from the sodium on the water burned 
with a yellow flame. The heat of this flame is intense, 
as is seen by the rapidity with which the glass jet becomes 
red hot. When hydrogen burns in the air or in oxygen, 
water is the only product, the union being expressed by 
the following equation : — 





or, 



2H tt + 







2 H 2 0. 




17. Combination of Hydrogen with Other Substances. — 

The explosiveness of hydrogen when mixed with oxygen 
has already been noticed. One of its most remarkable 
properties is that of being absorbed or occluded by certain 
metals. By finely divided platinum the absorption is so 
rapid that the metal becomes red hot, and the jet of 
hydrogen is quickly ignited. Likewise, if a piece of 
spongy platinum be lowered into a mixture of hydrogen 
and oxygen, the rapid absorption in a short time causes 
sufficient heat to explode the mixture. 

At the usual temperature, hydrogen has very little 



HYDROGEN 45 

affinity for most substances. As will be seen in Experi- 
ment 65, it explodes violently when mixed with chlorine, 
either on the approach of a strong light or by means of a 
spark. It unites vigorously with oxygen also on the 
application of a flame, but a light has no effect. 

18. Liquid Hydrogen. — Hydrogen is one of the most 
difficult gases to reduce to the liquid condition. This has 
been accomplished, however, by reducing the temperature 
to — 205° C, and allowing it to escape rapidly from a pres- 
sure of 180 atmospheres into a vacuum. At the same time 
this space is surrounded by a temperature of —200° C. 
Considerable quantities have been obtained in this way. 
In April of the year 1900, Dewar even succeeded in solidi- 
fying the gas. He surrounded liquid hydrogen with lique- 
fied air, and then by a pump caused so rapid an evaporation 
of the hydrogen that he soon obtained the remainder in a 
white, opaque solid. 

19. Uses of Hydrogen. — As a gas it has but few prac- 
tical uses. Its suitability for filling balloons has been 
mentioned, but in such cases it is generally used in a very 
impure form, mixed with various hydro-carbons given off 
with the hydrogen in the later distillation of coal. In the 
nascent state, that is, at the instant it is set free from some 
compound, it has great chemical activity and has the power 
of reducing many metals from their compounds. This 
use has already been seen in the passage of hydrogen over 
copper oxide, and will be further illustrated in our work 
with silver, iron, and other metals.* 

* The use of hydrogen in an*automatic cigar-lighter is occasionally seen. 
As shown in the figure, a small glass cylinder has a cubical block, a, of 
porcelain in the bottom : upon this rests an inverted glass cylinder, c, 
with a tubular neck and stop-cock, s ; above this jet is supported a plat- 
inum sponge, p ; under the small cylinder upon the porcelain block is 



46 



MODERN CHEMISTRY 



SUMMARY OF CHAPTER 

Hydrogen — Origin of term and meaning. 
Occurrence of hydrogen. 
Methods of making hydrogen. 
By decomposing water. 

With sodium or potassium. 
Describe method and apparatus. 
Chemical action — Proof of, by experiment. 
By decomposing acids. 
With zinc. 

Draw apparatus and explain method. 
Commercial methods. 
For filling balloons. 
Characteristics of hydrogen. 
Experiments to illustrate. 
Density. 
Inflamm ability. 
Explosiveness, etc. 
Liquid hydrogen. 
Uses of hydrogen. 
Special points. 

Explain the hydrogen pistol. 
Philosopher's lamp. 
Singing flame. 



placed some zinc, 0, and in the outer cylinder 
diluted sulphuric acid. As the acid and zinc 
react upon each other, hydrogen fills the in- 
verted cylinder, forces out the acid, and the 
action ceases. If now the stop-cock is opened, 
the hydrogen flows out of the jet, the acid 
reenters, and the generation of gas continues. 
As already seen, the hydrogen jet is quickly 
ignited by the platinum. When the customer 
has lighted his cigar, the stop-cock is again 
turned, and the action soon ceases. It ought 
to be said, perhaps, that such apparatus is 
of more interest as a novelty than as of real lasting utility. 







CHAPTER V 

OXYGEN, COMBUSTION, OZONE 

Oxygen : = 16 

1. Its Discovery. — The term oxygen is derived from 
two Greek words, meaning acid-former, and was given to 
this element because it was believed to be essential to 
the production of all acids. Oxygen was discovered by 
Scheele in 1773, but he did not publish his discovery until 
1775 ; and as in the meantime Priestley had isolated the 
same gas and had published an account of his experiments, 
the latter is generally given the credit. 

2. Abundance of Oxygen. — Oxygen is found in the 
atmosphere in large quantities, uncombined, constituting 
about one-fifth of the whole. It has been estimated that 
there is in the atmosphere alone over two and a half 
million billions of pounds. A liberal estimate of the 
amount used annually in respiration, and all forms of 
combustion, is about two and a quarter billion pounds. 
At this rate, in a century the entire world would use only 
one ten-thousandth part of the whole. At the same time 
it must be remembered that plant life is pouring the oxy- 
gen back again into the air, so that there is no danger of 
the equilibrium being destroyed. Oxygen also forms by 
weight eight-ninths of water, and being absorbed by the 
same, exists therein in considerable quantities in a free 
state. It is this free oxygen which is breathed by fishes. 
On account of its great affinity for other substances, it is 

47 



48 



MODERN CHEMISTRY 



found in combination with nearly all known elements, and 
forms in this way about 45 to 50 per cent of the earth's 
crust. 

3. How to produce Oxygen. — As a matter of historical 
interest, the method employed by Priestley is still some- 
times used. It is as follows : — 

Experiment 27. — Place in a hard-glass test-tube about a half 
gram of mercuric oxide, HgO, and heat strongly. Notice the change 
in the appearance of the oxide. Insert into the tube a pine splinter 
with a spark upon the end. What happens? What do you notice 
collecting upon the sides of the tube ? What substances have there- 
fore been obtained by heating this oxide ? 

4. Explanation. — The heat used has served to decom- 
pose the molecules of mercuric oxide into their constituent 
parts, thus : — 



+ HEAT — 





or, 



2 HgO + heat = 2 Hg + 2 . 



The two molecules of red oxide of mercury have yielded 
two molecules of mercury, one atom in each, and one 
molecule of oxygen, having two atoms. By continuing 
the operation the entire amount of the red oxide would 
disappear, while the deposit of mercury upon the sides of 
the tube would gradually increase. 

5. Other Methods of obtaining Oxygen. — The above 
method, though of interest, furnishes too limited a quan- 
tity of oxygen for practical purposes. A better and more 



fc 



with 




Fig. 12. 



OXYGEN, COMBUSTION, OZONE 49 

common way is to heat potassium chlorate, KCIO 
manganese dioxide, Mn0 2 . 

Experiment 28. — Mix together in a good-sized test-tube, or small 
flask, 1 or 2 g. of potassium chlorate and half as much manganese 
dioxide. Support upon a ring-stand 
with a wire screen protection, as shown 
in the accompanying figure, and attach 
the cork and delivery tube. Heat gen- 
tly at first, and then more strongly, but 
moderately, so as to regulate the flow 
of gas and not let it become too rapid. 
Allow the first that comes over to 
escape, then collect several bottles of 
the gas over water as you did the hydro- 
gen, and use for the following experi- 
ments. Save the contents of the flask 
for further use. 

Experiment 29. — Slip a sheet of glass or paper under a small 
bottle of oxygen, and place it in an upright position upon the table. 
Now plunge into the oxygen a taper, or pine splinter, with a spark 
upon it. Do you obtain the same results as before in the case of the 
oxide of mercury? 

Experiment 30. — Into another bottle of oxygen lower a deflagrat- 
ing spoon containing some burning sulphur; does it burn any differ- 
ently than in the air? If no deflagrating spoon is at hand, the student 
can prepare one by hollowing out the end of a short stick of gas 
carbon, or of crayon, and attaching a wire handle of suitable length. 

Experiment 31. — Fasten a piece of soft or bark charcoal to a 
stout iron wire, hold it in the burner flame until it begins to glow, 
then plunge into a jar of oxygen. If the charcoal is soft, the results 
will be very striking. Describe them. 

Experiment 32. — Twist together three or four fine iron wires, 
fasten to the end a small pine splinter, or warm and dip into sulphur 
and ignite. Plunge quickly into a large jar of oxygen which contains 
about an inch of water in the bottom. Describe the results. Do you 
see anything falling to the bottom of the bottle? A knife-blade or 
watch-spring may be thus burned, by first drawing the temper and 
using a larger amount of kindling material. 



50 



MODERN CHEMISTRY 



Experiment 33. — Put into a deflagrating spoon a small piece of 
phosphorus, ignite it, and thrust into a large jar of oxygen. Describe 
the combustion and the fumes that fill the jar. These are phosphorus 
pentoxide, a substance mentioned under deliquescent bodies. 

6. The Chemical Action. — In preparing oxygen as above 
from potassium chlorate and manganese dioxide, the latter 
remains unchanged at the close of the experiment. What 
\has really taken place, then, is the same as in the use of 
tKe mercuric oxide. The heat has simply decomposed the 
molecules of potassium chlorate, setting free the oxygen 
and \eaving behind a new compound containing only 
potassitipi and chlorine, called potassium chloride. The 
change may be shown thus : — 




+ HEAT = 




The two molecules of potassium chlorate shown here 
have each given up three atoms of oxygen, which have 
combined to form three molecules of oxygen, while two 
molecules of potassium chloride remain behind. These 
facts are more usually written thus : — 

2 KC10 3 + heat = 2 KC1 + 3 2 . 

7. Effect of the Manganese Dioxide. — If potassium 
chlorate be used alone, instead of mixing with manganese 
dioxide, as we did above, the same results are obtained, 



OXYGEN, COMBUSTION, OZONE 51 

but considerably more heat is required. Apparently, 
therefore, manganese dioxide has simply acted by its 
presence, or, as it is called, by catalysis. It is believed, 
however, that the dioxide is first converted into another 
compound, which at the temperature present is unstable, 
and that this in breaking up yields oxygen and the dioxide 
again. 

Experiment 34. — To prove that the manganese dioxide remains 
unchanged and that potassium chloride is formed. To the residue in 
the flask in Experiment 28, add about 50 cc. of water, let it stand a 
few minutes, shaking occasionally, warm gently and then filter. Boil 
this clear filtrate to dryness in an evaporating dish. While this is 
proceeding, add a little water to the black residue on the filter paper 
once or twice to wash it, throw the water away, and let the black 
residue dry. When the solution in the evaporating dish is perfectly 
dry, scrape it out, mix with a little fresh manganese dioxide, transfer 
to a test-tube, and heat. Do you observe any indication of oxygen 
being given off? If not, we may conclude under the present circum- 
stances that the oxygen was all removed in the previous heating, and 
that the white solid residue is KC1, potassium chloride, and not potas- 
sium chlorate, KC10 3 . 

When the black residue on the filter paper is dry, mix with it a 
little potassium chlorate, transfer to a test-tube and heat. Is oxygen 
given off readily? What proof? Is there any reason for believing 
that the black residue is still manganese dioxide ? 

8. The Proof by Weighing. — It will be well to try an 
experiment by which the facts discovered in preceding 
experiments may be proved. Such is the following experi- 
ment. 

Experiment 35. — Put into a test-tube or flask about a gram of 
potassium chlorate, put it upon the scales and balance it with shot or 
sand in a small box. Now add about a half grain of manganese 
dioxide, and then weigh carefully. As the box and shot counter- 
balance the flask and potassium chlorate, the weights added show at 
once the amount of dioxide used. Now connect a delivery tube and 
heat to drive off the oxygen. When the operation is complete, known 



52 MODERN CHEMISTRY 

by the fact that the gas no longer bubbles up through the water, re- 
move the delivery tube from the water, and let the flask cool. When 
cold, add a few cubic centimeters of water to dissolve the potassium 
chloride, then filter and wash the black residue as before. When 
thoroughly dry, weigh the residue and filter paper and subtract the 
weight of the paper. The latter may be obtained by weighing ten of 
them, or, if the balance is not very delicate, a hundred, and then tak- 
ing the fractional part. Does the weight of the black compound now 
agree with its weight before heating? 

9. Commercial Methods of making Oxygen. — Most of 
these methods consist in abstracting oxygen from the air 
by using a substance which when heated or when under 
pressure will absorb oxygen, and then w^hen cooled or 
when the pressure is removed will again give it up. One 
of the best known of these methods is Brin's, which con- 
sists in using barium oxide, BaO, as the chemical agent. 
When gently heated in the air, it takes up an additional 
amount of oxygen, forming barium dioxide, Ba0 2 , thus, — 

BaO + O = Ba0 2 . 

If, now, the heat is increased, the barium dioxide is unable 
to retain the additional atom of oxygen taken from the 
air and gives it up again, thus, — 

Ba0 2 + heat = BaO + O. 

Or, if the pressure under which the barium dioxide was 
formed is decreased, the same results follow at considera- 
bly less expense. 

10. Motay's Method. — The principle of this is about 
the same as that of Brin's, but different substances are 
used. Manganese dioxide and caustic soda, when heated 
moderately in a current of air, form a compound which 
at a higher temperature is again decomposed, yielding up 
the oxygen previously taken from the air. 



OXYGEN, COMBUSTION, OZONE 



53 



11. Other Methods. — These are not important as a 
means of producing oxygen for commercial or experimental 
purposes, but the principle underlying them is involved 
in a number of the processes of chemistry with which we 
shall deal later, and should consequently be understood. 
It will be noticed that in all the methods of preparing 
oxygen used above, we have employed substances con- 
taining a large per cent of that element. There are sev- 
eral other substances of similar composition which may 
be made to furnish oxygen. Thus, manganese dioxide, 
Mn0 2 , potassium dichromate, K 2 Cr 2 7 , and potassium 
permanganate, KMn0 4 , when heated with sulphuric acid, 
H 2 S0 4 , will yield oxygen. 

Experiment 36. — Put a half gram of manganese dioxide into a 
test-tube and add about a cubic centimeter of sulphuric acid. Warm 
gently, collect a small bottle of the gas, and make the usual test for 
oxygen. What are the results ? 

It will be found that the dichromate and the permanga- 
nate act in a similar way, except that the quantity of gas 
obtained is considerably greater. The chemical action 
may be shown thus : — 

Mn0 9 + H 9 S0. = MnS0 4 + H 9 + O. 




54 MODERN CHEMISTRY 

In a similar way the reaction of potassium dichromate 
and sulphuric acid upon each other may be shown ; 

K 2 Cr 2 7 + 4 H 2 S0 4 

= K 2 S0 4 + Cr 2 (S0 4 ) 3 + 4 H 2 + 3 O, 

and of potassium permanganate and sulphuric acid, 



2 KMn0 4 + 3 H 2 S0 4 



K 2 S0 4 + 2 MnS0 4 + 3 H 9 + 5 O. 



See pages 322, 325, for application of this property of the 
above substances. 

12. Characteristics of Oxygen. — Oxygen is an odorless, 
colorless gas, a little heavier than air, the weight of a liter 
being 1.43 g. As already noted it is slightly soluble in 
water, and upon this fact depends the life of aquatic ani- 
mals, which abstract this free oxygen from the water. It 
may be liquefied by extreme cold and pressure. This was 
first accomplished about a quarter of a century ago by 
Cailletet and Pictet, who succeeded in preparing a small 
quantity at great cost. At present it is made in almost 
any amount by first liquefying air and then allowing the 
nitrogen to boil out. (See page 100.) 

13. Peculiarities of Liquid Oxygen. — In the liquid con- 
dition oxygen is of a pale blue color and boils at about 
— 180° C, a few degrees higher than the boiling point of 
nitrogen. It presents many striking peculiarities ; a rod 
of carbon heated red hot and plunged into the liquid 
oxygen at a temperature 180° below zero burns vigorously, 
while a stout iron wire similarly heated is rapidly con- 
sumed with a brilliant display of sparks. Cotton rags 
saturated with it and confined in a cylinder, when ignited, 



OXYGEN, COMBUSTION, OZONE 55 

explode so violently as to burst tubes made of iron or 
brass. 

14. Chemical Affinity of Oxygen. — The strongest chem- 
ical property of oxygen is its affinity for other substances. 
This was seen in the rapidity of combustion of the various 
ignited substances when placed in an atmosphere of oxy- 
gen. From these experiments it is not difficult to see 
what would be the results were the air undiluted oxygen. 
The smallest spark would be sufficient to start the fiercest 
conflagration, while our stoves, furnaces, etc., would be 
rapidly consumed, accompanied by a most brilliant display 
of sparks. 

15. Uses of Oxygen. — As is well known, oxygen is 
absolutely necessary for life. It is absorbed by the blood 
through the walls of the air-cells of the lungs and carried 
by the red corpuscles to all parts of the body. Here it 
unites with the waste material, burning it to carbon 
dioxide and other compounds, and at the same time warm- 
ing the body. The carbon dioxide is carried back to the 
lungs, from which it is thrown off into the air. In cases 
of asphyxiation pure oxygen is sometimes used as a restora- 
tive, but ordinarily, if breathed for any length of time, the 
temperature of the body rises owing to the increased de- 
struction and consumption of tissue, and general feverish 
symptoms follow. A limited number of experiments by 
the author show that small animals, such as mice, when 
placed in an atmosphere of pure oxygen soon exhibit great 
activity, followed by apparent relaxation and complete 
exhaustion. Experiments have also shown that animals 
in oxygen under pressure would very quickly die, as if the 
gas in this condition were an active poison. 

Experiment 37. — Into a hard-glass tube, T, the weight of which 
is known, supported in a ring-stand as shown in Fig. 13, put 2.5 g. 



56 



MODERN CHEMISTRY 



of potassium chlorate. Let the tube, f, just reach through the cork 
of A y while the tube, c, extends nearly to the bottom. B needs no 

cork, but those in A and T 
should be of rubber and fit air- 
tight. Heat the contents of 
the tube gradually and then 
strongly until the water is no 
longer forced from A into B. 
Allow the tube to cool, and 
w r eigh. The loss wdll represent 
the oxygen expelled. Measure 
carefully the water in B. This 
will represent the volume of 
oxygen at the pressure and 
temperature of the room. For 




Apparatus to use in finding 
the Weight of Oxygen. 



methods of reducing this to standard conditions, see page 96. 
Suppose the tube weighs 15 g. ; then 



tube + KC10 3 

tube + KC1 

17.5 - m 



-. 17.5 g., before heating ; 
= m grams, after heating ; 
= wt. of oxygen expelled. 



Volume of H 2 in B = n cc, or what is the same, n cc. of oxygen 
at existing temperature and pressure, and by calculation, n 1 , at stand- 
ard conditions. Then ?i' cc. weigh (17.5 — m) grams, from which the 
student can easily find the weight of one liter or 1000 cc. This should 
be 1.43 g. Let the student determine what is his per cent of error. 

By similar methods, the weight of a liter of various 
other gases, insoluble in water, may also be determined. 

16. Oxidation and Combustion. — When any substance 
combines with oxygen to form a new compound, it is said 
to be oxidized, and the process is known as oxidation. 
This may be slow or rapid. When it takes place so 
rapidly as to be accompanied by heat and light, the pro- 
cess is called combustion. To illustrate: when a piece of 
iron is exposed to the air in the presence of moisture, it 
soon becomes covered with rust, which is really an oxide of 



OXYGEN, COMBUSTION, OZONE 



57 



iron; in other words, the iron has been oxidized. Again, 
when we tipped the iron wires with sulphur and ignited 
it, they were rapidly consumed in the jar of oxygen with 
much heat and considerable light. This was combustion. 
A pile of brush will gradually decay, or oxidize, without 
any perceptible heat, but by setting it on fire we quickly 
destroy it by the process of combustion. 

17. Combustible Substances and Supporters of Combus- 
tion. — Substances which thus burn in oxygen or its 
diluted form, the air, are said to be combustible, while the 
substance in which they burn is called a supporter of com- 
bustion. Thus, when a jet of hydrogen burns in a jar of 
oxygen, the former would be 
spoken of as the combustible 
substance, and the latter as 
the supporter of combustion. 
It is true, however, that if we 
thrust a delivery tube from 
which a current of oxygen is 
issuing, up into a jar of 
hydrogen which is burning at 
the mouth, as seen in Fig. 14, 
the jet of oxygen will be seen to burn in the atmosphere 
of hydrogen, just as before the hydrogen did in the oxy- 
gen. Yet in view of all the facts it seems better to adhere 
to the statement previously made, that it is really the 
hydrogen which burns, and the oxygen which supports 
the combustion. 

18. Kindling Temperature. — It is well known that some 
substances ignite much more readily than others. This, 
chemically speaking, simply means that some combine 
with oxygen at a lower temperature, or much more readily, 
than do others. Thus, substances like alcohol and many 




Fig. 14. 



58 MODERN CHEMISTRY 

oils need but little heat to ignite them; phosphorus, like- 
wise. Pine wood needs a higher temperature, and hard 
wood still higher. The point at which any substance 
takes fire is said to be its kindling temperature. 

19. What is a Flame? — A flame is simply burning gas. 
Whenever a substance will not burn with a flame, it is 
because there is either no gas present or there is nothing 
which may be converted into a gas. For example, when 
a lamp burns, the oil drawn up through the wick by capil- 
lary attraction is volatilized by the heat, and it is the 
burning of this gas that makes the flame. On the other 
hand, charcoal and the hardest natural coals do not burn 
with a flame, because previous heating has driven out all 
the gaseous products. However, they may be heated suf- 
ficiently to be partially converted into carbon monoxide, 
a gas which burns with a pale blue flame. 

20. The Oxy-hydrogen Blowpipe. — This is a lamp 
arranged for burning hydrogen thoroughly mixed with 
oxygen, and affords one of the hottest flames known. 

Its construction will be understood 
from the figure, which gives a 
sectional view. The inner tube, 
m, is connected with the oxygen 
tank, and the outer, n, with the 
hydrogen. In this way, as the 
inner tube is somewhat shorter, 
the gases become thoroughly mixed 
before leaving the tube at E, hence 
the combustion is perfect. The 
Fig. 15. - The Oxy-hydro- pressure should be so regulated 

gen Blowpipe. . . ° 

as to furnish twice as much hydro- 
gen by volume as oxygen. This blowpipe is used for 
melting very refractory substances. It is also used espe- 




OXYGEN, COMBUSTION, OZONE 59 

cially in furnishing light for stage effects, stereopticon 
views, and for illuminating moving floats in street parades 
given after dark. When used for these purposes the 
almost non-luminous blue flame is allowed to strike upon 
a stick of prepared lime supported in a socket just in 
front of the blowpipe. This is often called the calcium 
or Drummond light, and is of dazzling whiteness, rivaling 
the electric arc light. 

Ozone 

21. Its Discovery. — This substance, on account of its 
peculiar odor, was named from a Greek word which means 
to smell. It was first observed in passing electrical sparks 
through a tube of oxygen, and is always noticeable when 
an electric discharge takes place in the air. 

22. What is Ozone ? — For some time it was regarded as 
a compound body, but is now known to be simply a con- 
densed form of oxygen. Quite a number of substances, 
such as sulphur and phosphorus, appear in a form other 
than the usual one : this is known as the allotropic, a word 
which means simply another form. 

23. Methods of obtaining Ozone. — It is impossible to 
prepare ozone in large quantities in a pure condition, be- 
cause by the best methods only a small per cent of the 
oxygen used is converted into its allotropic form. Usu- 
ally not over one or two per cent is obtained, and under 
the most favorable circumstances only about twenty per 
cent. In the ordinary methods of making oxygen, ozone 
is almost always obtained in appreciable quantities. One 
of the easiest methods of preparing it is given in the fol- 
lowing : — 

Experiment 38. — Scrape a stick of phosphorus perfectly clean, 
put it into a bottle, and add water sufficient to cover about half of it. 
In a few minutes the presence of ozone may be detected by suspending 



60 MODERN CHEMISTRY 

in the bottle a strip of white paper moistened with a solution of 
potassium iodide and starch. The paper will turn decidedly blue. A 
solution of potassium permanganate treated with strong sulphuric 
acid also gives the test for ozone, along with the oxygen thus 
evolved. 

24. Ozone in the Air. — Ozone is believed to exist in 
appreciable quantities in the atmosphere, being produced 
mainly by electrical discharges. Its presence in dwell- 
ings is never perceptible, and scarcely ever in large cities, 
but in the country a strip of starch paper exposed to the 
breeze for some time shows the characteristic blue color. 

25. How Ozone differs from Oxygen. — As already stated, 
it is a condensed form of oxygen. Experiment has shown 
that if a given amount of ozone is decomposed so as to 
form ordinary oxygen, the volume increases one-half ; that 
is, 100 cc. of ozone would become 150 cc. of oxygen. On 
the other hand, if a closed volume of oxygen be subjected 
to a silent discharge of electricity so as to convert a por- 
tion of it into ozone, a corresponding decrease of volume 
takes place. 

26. For example, suppose 150 cc. of oxygen be thus 
treated, and that 30 cc. are converted into ozone. It is 

found that by absorbing this ozone 
so as to separate it from the remain- 
ing oxygen, and again setting it 
free, there are only 20 cc. of ozone, 
Molecule of Molecule of w hile but 120 cc. of oxygen remain. 

Oxygen. Ozone. ° ° 

If, however, this 20 cc. of ozone be 
heated strongly, so as to convert it into oxygen again, 
we shall find the volume increases to 30 cc. The mole- 
cule of ozone therefore would differ from that of oxygen, 
in that it contains three atoms of oxygen, while the other 
has only two. This is shown in the accompanying figure. 





OXYGEN, COMBUSTION, OZONE 61 

From this it naturally follows that ozone is 50% more 
dense than oxygen. 

27. Properties of Ozone. — The properties are also dif- 
ferent from those of oxygen. Its odor has already been 
mentioned. If placed in a long glass tube, so as to give 
considerable depth, it is seen to have a blue tinge. It 
readily destroys the color of such vegetable solutions as 
indigo and litmus and quickly attacks such metals as mer- 
cury and silver, which remain unchanged in the air and 
which are even unaffected by oxygen when heated. We 
have seen its effect upon potassium iodide above. Free 
iodine always turns starch blue. In the test made, the 
ozone united with the potassium in the potassium iodide to 
form an oxide with the metal, and the iodine was thus set 
free. In the same way if a drop of ammonium hydroxide 
be let fall into a jar of ozone, a dense white cloud forms, 
owing to the fact that a white solid compound of ammonia 
is formed, thus : — 

2 NH 4 OH + 3 = NH 4 N0 2 + 3 H 2 0. 

28. Liquid Ozone. — Ozone may be liquefied at ordinary 
atmospheric pressure by reducing the temperature to 
— 106° C, — a point considerably higher than that at which 

.oxygen liquefies. Ozone is also a very unstable body, 
changing back readily into oxygen ; an illustration of this 
is seen in the fact that if a quantity of ozone be suddenly 
compressed and heated, it explodes with violence. It is 
because of this instability that ozone is so strong an oxi- 
dizing agent. The nascent oxygen liberated, if inhaled, 
attacks the mucous linings, causing an irritation some- 
what like that of dilute chlorine. More than this, head- 
ache soon follows, if much ozone is inhaled, even though 
diluted with considerable quantities of oxygen. 



62 



MODERN CHEMISTRY 



TABULAR VIEW OF DIFFERENCES 



Oxygen 


Ozone 


Colorless. 


Blue. 


Odorless. 


Peculiar odor. 


Density ; slightly heavier 


Density ; considerably 


than air. 


heavier than air. 


Two atoms in molecule. 


Three atoms in molecule. 


Strong oxidizer. 


Very strong oxidizer. 


Liquefies at - 180° C. 


Liquefies at - 106° C. 


Stable. 


Unstable. 



29. Uses of Ozone. — It is believed to have a beneficial 
effect in destroying disease germs and in oxidizing decay- 
ing organic matter. 

30. Isomeric and Polymeric Bodies. — Just as ozone is 
another form of oxygen, so we shall find that phosphorus 
and certain other elements present allotropic forms as 
unlike the usual forms as oxygen and ozone. When we 
come to the study of compound bodies we often find 
two substances not at all alike in properties, which, upon 
analysis, are found to contain exactly the same elements 
united in exactly the same ratios. Thus aldehyde and 
oxide of ethylene both have the same composition, repre- 
sented by the formula C 2 H 4 0, but their properties are 
very different. Such substances are said to be isomeric. 

31. Sometimes while they have the same percentage 
composition, the vapor density of one will be several 
times that of the other. Thus acetylene is C 2 H 2 , and 
benzine C 6 H 6 . In each case the carbon is If, or 92.3 per 
cent, of the molecule ; but the molecular weight of one 



OXYGEN, COMBUSTION, OZONE 63 

is three times that of the other. Such substances are 
said to be 'polymeric. 

Hydrogen Dioxide: H 2 2 

32. Composition. — This is a compound which in some 
of its characteristics resembles ozone. In composition it 
is most like water, having one additional atom of oxygen. 
It is believed to exist in the air in minute quantities, and 
some of the effects attributed to ozone may be due to 
hydrogen dioxide. 

33. How to obtain It. — For experimental purposes it is 
usually prepared by treating barium dioxide with dilute 
sulphuric or hydrochloric acid. 

Experiment 39. — Add to about a gram of barium dioxide a little 
water, and then dilute sulphuric or hydrochloric acid. Stir for a 
moment or two with a glass rod. 

To prove the presence of hydrogen dioxide, add a few drops of 
potassium dichromate and about a half cubic centimeter of ether, and 
shake well. The hydrogen dioxide forms a blue solution with the 
dichromate, which is taken up by the ether and thus concentrated 
within little space. 

34. Some of its Peculiarities. — Like water, it is a color- 
less liquid, but is thicker or sirupy, and has a bitter 
taste. It is very unstable, decomposing at all tempera- 
tures into water and oxygen ; it is therefore a good 
bleaching agent, the bleaching being done by the nascent 
oxygen. Like ozone, it readily tarnishes silver and decom- 
poses potassium iodide, giving in the same way a test 
with starch paper. It is soluble in water, and thus 
diluted it will bleach the skin, but when concentrated it 
burns or blisters it. 

35. Uses. — This compound, more usually sold under 
the name hydrogen peroxide, is now manufactured very 



64 MODERN CHEMISTBY 

cheaply, and is used to a considerable extent as a bleach- 
ing agent, especially for hair and feathers. It is used 
largely by dentists and in surgery as an antiseptic, and to 
some extent in cleansing oil paintings and engravings. 

SUMMARY OF CHAPTER 

Discovery of oxygen. 

Meaning of term. 
Abundance of oxygen — Various forms in which it occurs. 
Methods of preparing oxygen. 
Priestley's method. 
Ordinary method. 

Chemicals and apparatus used. 
Chemical changes involved. 

Proof of these changes — Experimental. 
Proof by weight. 
Other methods. 

Character of substances used. 
Characteristics of oxygen — Compare with hydrogen in 
Color, odor, density, combustibility. 
Power of supporting combustion. 
How would you distinguish the two ? 
Some peculiarities of liquid oxygen. 
Uses of oxygen. 
Special. 

Meaning of terms combustion, oxidation, flame, kindling point. 

Illustration of the terms. 

Description and drawing of oxyhydrogen blowpipe. 

Uses for it. 
Ozone — Meaning of the word. 

What is its relation to oxygen ? 
Methods of obtaining. 

Compare with oxygen, showing differences. 
Value of ozone. 
Hydrogen dioxide — Formula. 

Compare with water in properties. 
Uses for the compound. 



CHAPTER VI 

CHEMICAL NOTATION, SYMBOLS, FORMULA, EQUATIONS, 

PROBLEMS 

1. Symbols. — The student will have noticed that in 
chemistry we frequently employ a short-hand method of 
expressing the different elements and their compounds. 
Thus we have seen that hydrogen is represented by H, 
oxygen by O, and so on. These are called symbols. 

2. Their Form. — Frequently, the symbol of an element 
is its initial letter ; often, however, this is the same for a 
number of elements, as for example, carbon, calcium, cad- 
mium, copper, etc. In such cases, the most common 
usually is designated by the initial letter; another by the 
first and second letter, as Ca, calcium; another by the first 
and some other distinctive letter, as Cd for cadmium. 
Frequently, the first or the first and second letters of the 
Latin term for the same substance are used, as Cu for 
copper, from cuprum. In like manner, sulphur, silicon, 
selenium, silver, are designated by the symbols S, Si, Se, 
and Ag (from the Latin argentwii). The Latin has * fur- 
nished a number of the symbols of the common elements : 
thus, sodium, Na (natrium), potassium, K (kaliurn)^ iron, 
Fe (ferrum). The sj^mbol, Hg (hydrargyrum), for mer- 
cury is from the Greek. 

3. Strict Meaning. — Strictly speaking, the symbol of 
an element not only represents that element, but a defi- 
nite amount of it : that is, one atom. Hence, to speak of 
an element by using its symbol when we mean an indefinite 
amount is unscientific and should not be practiced. 

65 



66 MODERN CHEMISTRY 

4. Formulae. — As the elements are represented by- 
symbols, so compound bodies are by formulae ; that is, by 
an aggregation of symbols. Compounds are usually 
named by simply combining the terms representing the 
elements entering into the composition, the more electro- 
positive being placed first ; thus, potassium iodide consists 
of two elements, potassium and iodine. The formulae are 
always arranged in the same way: thus, KI, potassium 
iodide ; KC1, potassium chloride. It will be seen, there- 
fore, that as a symbol represents an atom of an element, 
so a formula represents the smallest amount of a compound 
body, a molecule. 

5. Sub-figures. — When the elements enter into com- 
position, in other than a single atom of each, that fact is 
indicated by putting a small figure below and at the right 
of the symbol ; thus, H 2 0, the formula for water, indi- 
cates that there are two atoms of hydrogen in the mole- 
cule, and, H 2 2 , for hydrogen peroxide, indicates that 
there are two atoms of each. These sub-figures are some- 
times appended to a group of elements, in which case the 
group is inclosed in parentheses and the figure placed 
outside ; for example, lime-water, or calcium hydroxide is 
Ca(OH) 2 . This might also be written Ca0 2 H 2 , but the 
former method is preferable, as will be seen later. If we 
desire to indicate more than one molecule of a substance, 
this is done by prefixing a coefficient to the formula. 
Thus, 2 HC1 indicates two molecules of hydrochloric acid ; 
5 H 2 0, five molecules of water. 

6. Radicals. — By a radical we mean a group of ele- 
ments which in most chemical reactions seem to hold 
together, but which do not by themselves form a distinc- 
tive compound. For example, (HO) seen in the formula 
for lime-water is a radical known as hydroxyl, which enters 



REACTIONS 67 

into a great many compounds. Again, (NH 4 ) is a group 
called ammonium, which is very common, and ordinary 
ammonia water is NH 4 OH, composed of two radicals 
(XH 4 ) and (OH), not written NH 5 0, because of this fact. 

7. Reactions. — Equations in chemistry which show 
the chemical changes that take place when two or more 
substances react with each other are called reactions. We 
have already seen a number of these ; thus : — 

2 Na + 2 H 2 = 2 NaOH + H 2 ; 
Zn + H 2 S0 4 = ZnS0 4 + H 2 . 

8. The first indicates that two atoms of sodium uniting 
with two molecules of water will produce two molecules 
of caustic soda and two atoms or one molecule of hvdro- 
gen. Another thing must be noticed, and that is that 
every atom appearing in one member of the equation must 
also be found in the other. Thus, the two atoms of 
sodium are seen in the second member of the equation in 
the two molecules of caustic soda, the four atoms of hydro- 
gen in the water appear partly in the caustic soda and 
partly as free hydrogen ; likewise the two atoms of oxy- 
gen in the water are found in the caustic soda. It must 
be borne in mind that a coefficient before a formula mul- 
tiplies every symbol in that formula. Thus, 
2 KClOg means that there are two atoms of 
potassium, K; two of chlorine, CI; and six 
of oxygen. This will be seen from the fol- 
lowing illustration : — 

a represents a molecule of water containing two 
atoms of hydrogen and one of oxygen ; b represents 
a second molecule having the same composition. Tak- 
ing both together, or two molecules of water, 2 H 2 0, we see there are 
four atoms of hydrogen and two of oxygen. 




68 MODERN CHEMISTRY 

9. Atomic Weights. — We cannot think of matter with- 
out assigning to it some weight. So the atoms of the 
elements, though the smallest conceivable portions of 
matter, are assumed to have definite weights. Hydrogen, 
being the lightest of substances, is taken as the standard,* 
and its atomic weight is assumed to be one, or by some, 
as one micro-crith. Of course, a weight as small as this 
has never been determined, and is therefore merely an 
abstract idea; but something is necessary for comparison. 
When we speak of the atomic weight of an element, there- 
fore, we simply mean its density compared with hydro- 
gen. Thus, we say the atomic weight of oxygen is 16, of 
carbon, 12 ; we mean that these elements are, respectively, 
sixteen and twelve times as heavy as hydrogen, or, if one 
cubic foot of hydrogen weighs a gram, one cubic foot of 
oxygen will weigh sixteen grams. 

10. Molecular Weight. — By molecular weight we mean 
the sum of the weights of the atoms entering into the 
composition of the molecule. For example, H 2 repre 
sents a molecule of water ; the two atoms of hydrogen 
weigh 2, the one of oxygen, 16, or all together, 18. The 
molecular weight of water is therefore 18. Now if we 
examine any chemical equation, we will find that the sum 
of atomic weights in one member must equal the sum of 
those in the other member. Take the following : — 

Na + H 2 == NaOH + H, 

and substituting the atomic weights as given in the table 
on page 9, we have 

23 + (2 + 16) = (23 + 16 + 1) + 1, or 41 = 41 ; 
and this must always be so in any true reaction. 

* There has long been a controversy whether Hydrogen, H = 1, or 
Oxygen, = 16, should be the standard. The latter is increasing in favor. 



EQUATIONS 69 

11. Writing Equations. — A chemical equation is valu- 
able in that it shows at once in concise form not only the 
substances which enter into the reaction, but also the 
products formed, and the exact amount of each. At first 
the student will experience some difficulty in completing 
even the simpler reactions, but it soon becomes a very 
easy matter. In writing the second member, always put 
down first the product or products that you know are 
produced, and then note what remains. There will 
usually form of the remaining elements some compound 
that is familiar, so that the ordinary reaction becomes 
very simple. Thus, suppose we are preparing ammonia 
from ammonium chloride, NH 4 C1, and caustic soda, 
NaOH. 

r NH 3 (product we know is produced) 
remaining H, Na, CI, H, O, from 
which water, H 2 and NaCl, form. 



NH 4 G1 + NaOH 



12. Practical Value of the Equation. — Having deter- 
mined by experiment the products that are formed in any 
chemical reaction, and having therefrom written the equa- 
tion, we can readily ascertain the amount of each product 
that will be formed from a certain amount of another ; or 
if required to produce a definite quantity of any body, we 
can calculate what it will be necessary to use in obtaining 
it. This is very important to the manufacturer, and only 
by paying the closest attention to this point can he hope 
to succeed. To illustrate, suppose we are required to 
determine how much zinc will be necessary for the prepa- 
ration of 50 g. of hydrogen. We would first write the 
equation, showing the preparation of hydrogen : — 

Zn + H 2 S0 4 = H 2 + ZnS0 4 . 



70 MODERN CHEMISTRY 

In the table we find the atomic weight of zinc is 65 ; 
looking at the reaction, then, we would see that 65 parts 
of zinc by weight produce, when reacting with the acid, 
2 parts of hydrogen. The gas obtained is therefore ^ by 
weight of the metal used. 

Then, 50 g. = &. 

^ = i of 50 g. = 25 g. 1 

|f = 65 x 25 g. = 1625 g. of Zn. 

Problem 1. — How much sulphuric acid is it necessary to put with 
260 g. of zinc in preparing hydrogen ? 

Using the same reaction as above, we find first the molecular weight 
of H 2 S0 4 , which is 98. We see then that 65 parts of zinc unite with 
98 of acid, or the acid used is |f of the metal. 

Then, || of 260 g. Zn = 98 X ^ 26Q = 392 g. H 2 S0 4 . 

65 

Some prefer to solve such problems by proportion, thus : — 

The wt. of the Zn : wt. of acid : : wt. of Zn in g. : wt. of acid in g. ; 
or, 65 : 98 : : 260 : x. 

g= 98x260 = 392> 
65 

Problem 2. — In using 260 g. of zinc in preparing hydrogen, how 
much zinc sulphate, ZnS0 4 , will be obtained ? 

Problem 3. — How many grams of oxygen may be obtained from 
450 grams of potassium chlorate? 

Problem 4. — How much caustic soda will be produced in prepar- 
ing 10 g. of hydrogen by using metallic sodium and water? 



SUMMARY OF CHAPTER 

Symbols and form ulse — Difference between them. 
Composed of what. 
Exact meaning of each. 



NITROGEN AND ITS COMPOUNDS 71 

Radicals. Atomic and molecular weights. 

Meaning of the term. Meaning of the terms. 

Illustrations. Illustrations. 

Chemical equations. 
Value of. 
Problems. 



CHAPTER VII 
NITROGEN AND ITS COMPOUNDS 

Nitrogen : N = 14 

1. History. — Nitrogen, meaning niter producer, was 
given this name because of its being an important con- 
stituent of saltpeter, often called niter. It had previously 
been called azote, a term which meant that it would not 
support life. 

2. Where found. — As already stated, nitrogen con- 
stitutes about four-fifths of the air, and is uncombined. 
It also exists in various compounds, such as saltpeter, 
potassium nitrate, KN0 3 , and Chile saltpeter, sodium 
nitrate, NaN0 3 . It also enters into the composition of 
many vegetable and animal products, and in their decom- 
position is given off into the air in the form of ammonia, 
NH 3 . 

3. How to prepare Nitrogen. — As nitrogen exists so 
abundantly in a free state in the air, this is the best source 
from which to obtain it. Any method by which we can 
remove the oxygen and leave the nitrogen will do. For 
this purpose phosphorus is generally used. 

Experiment 40. — Cover a large flat cork with a coating of plaster 
of paris and float it upon a pan or basin of water. A small iron 
saucer serves well instead of the cork. Put upon it a small piece of 



72 



MODERN CHEMISTRY 



phosphorus and ignite. Quickly place over the burning phosphorus 
a large wide-mouthed jar. Notice that the water gradually rises in 
the jar to take the place of the consumed oxygen, and that in a few 
minutes the white fumes are absorbed by the water. Owing to the 
expansion caused by the heat some bubbles of air almost always 
escape in the early part of the experiment. 

If the phosphorus is not ignited, but the combination with the 
oxygen is allowed to take place slowly, 
this loss may be avoided, but several hours 
are required. Sometimes it is more con- 
venient to use a deflagrating spoon instead 
of a cork ; if so, the handle must be bent 
V-shaped, so as to bring the phosphorus 
above the water even after it has risen in 
the jar. Notice about how much the 
water rises. When the fumes have all 
j,-, disappeared, lift the jar and put a burning 

candle up into the gas. What happens? 
Compare it with similar tests with oxygen and with hydrogen. 




4. Other Ways of preparing Nitrogen. — The method 
already given, while the easiest and most commonly used, 
does not give as pure nitrogen as may be obtained in 
some other ways. If a current of air be made to stream 
slowly over a tube con- 

Alr r 



taming copper turnings 
heated to redness, the oxy- 
gen will combine with the 
copper, forming copper ox- 
ide, and the nitrogen will 
remain. Then if this is 

allowed to bubble through Fig. 17. — Nitrogen, prepared bypass- 

a bottle of lime-water, the ing Alr over Copper - 

carbon dioxide will be absorbed, and we shall obtain a 
fairly pure nitrogen. The illustration will show the 
method. 




NITROGEN AND ITS COMPOUNDS 73 

Nitrogen may also be obtained by heating certain com- 
pounds containing it. 

Experiment 41. — Into a small flask put 1 or 2 g. of sal am- 
moniac, XH 4 C1, and the same amount of sodium nitrite, and add 
about 30 cc. of water. Heat gently and cautiously, and collect the 
gas over water as you did oxygen and hydrogen. Test the gas for 
nitrogen. What are your conclusions? 

5. Peculiarities of Nitrogen. — From the experiments 
made the student will notice that the gas has no color ; 
it is odorless, lighter than air, will neither burn as does 
hydrogen, nor support combustion as does oxygen. It 
has no affinity for other substances at ordinary tempera- 
tures. It will combine with red-hot magnesium in the 
absence of oxygen, and with oxygen when a discharge 
of electricity takes place, both of which methods have 
been used in preparing argon from its mixture with atmos- 
pheric nitrogen. It will not support respiration any 
more than it will combustion, and is one of the most inac- 
tive substances known. This inactivity, or feeble chemi- 
cal affinity of nitrogen, is the reason for the instability of 
man)' of its compounds, as seen in the explosiveness of 
gunpowder and nitroglycerine. 

6. Use of Nitrogen. — The use of nitrogen, except in 
the form of many valuable compounds, seems to be simply 
to dilute the oxygen of the air as already stated. 

Compounds of Nitrogen 

7. In an indirect way nitrogen forms a large number 
of compounds- many of which are very valuable. Among 
these we shall first consider ammonia. 

8. Ammonia, NH 3 . — As already stated, ammonia is one 
of the products formed in the decomposition of nitroge- 



74 MODERN CHEMISTRY 






nous organic matter ; that is, organic matter which con- 
tains nitrogen in addition to the usual carbon, hydrogen, 
and oxygen. It finds its way into the air from these 
sources, and being absorbed by the moisture of the air is 
brought down in the rain, and usually exists in very small 
quantities in cistern and river water. With these excep- 
tions, ammonia does not occur free to any extent, but is 
found abundantly in certain compounds, especially sal 
ammoniac or ammonium chloride, NH 4 C1. 

9. The commercial supply of ammonia is obtained from 
the distillation of coal in the manufacture of common 
\lluminating gas. (Seepage 153.) The decay of organic 
matter, attended by the formation of ammonia, occurs as 
follows : when the nitrogenous matter decomposes, the 
former arrangement existing among the atoms of carbon, 
oxygen, nitrogen, and hydrogen is broken up, and in the 
rearrangement the nitrogen and hydrogen unite to form 
ammonia. 

10. Ammonia prepared from Coal. — In the distillation 
of coal the process is really the same, but more rapid, and 
ammonia is one of the impurities given off with the hydro- 
carbon gases. These are all passed through a tank filled 
with water, which absorbs the ammonia and forms an 
aqueous solution, known as aqua ammonia or ammonium 
hydroxide. This, more or less impure, is drawn off at inter- 
vals and treated with hydrochloric acid, which converts it 
into a salt of ammonia, ammonium chloride, NH 4 C1, as 
shown by the following reaction : — 

NH 4 OH + HC1 = NH 4 C1 + H 2 0. 

11. Then by treating this chloride with some strong 
alkali like caustic potash or soda, and heating, ammonia 
is again liberated, and being passed into water, produces 



. 



NITROGEN AND ITS COMPOUNDS 



75 



the aqua ammonia of commerce. The following shows 
the reaction which takes place : — 

NH 4 C1 + KOH = NH 3 + KC1 + H 2 0. 

12. On account of its cheapness, slaked lime, Ca(OH) 2 , 
a compound very similar in properties to caustic potash or 
soda, is ordinarily used with the sal ammoniac to liberate 
the ammonia. The reaction is seen below : — 

2 NH 4 C1 + Ca(OH) 2 = 2 NH 3 + CaCl 2 + 2 H 2 0. 




Fig. 18. — Preparation of Ammonia. 

m, n, o, cylinders containing solutions of impure ammonium chlo- 
ride, as obtained from coal-gas factories, mixed with lime ; S, S, S, 
stirrers to keep the lime from settling ; F, furnace to heat the mixture 
and expel the ammonia ; P, B, condensers for cooling the ammonia 
gas ; C, cylinder of pure water to absorb the ammonia and thus form 
aqua ammonia; 7), trough of acid to combine with any fumes escap- 
ing from C- In this trough, if hydrochloric acid is used, there would 
form ammonium chloride. 

Experiment 42. — To illustrate the preparation of ammonia. Put 
about a half gram of sal ammoniac, XH 4 C1, into a test-tube and add 
to it about 1 cc. of water, then a little caustic soda or potash solution, 



76 MODERN CHEMISTRY 

and heat gently. Is there any gas given off: having an odor ? Hold 
in the mouth of the test-tube a piece of moistened red litmus paper 
and note the effects. Try also a piece of turmeric paper in the same 
way. How is it affected ? 

Experiment 43. — To about 2 g. of ammonium chloride in a 
tube or flask add 1 or 2 cc. of slaked lime, made by adding a little 
water to some lime ; adjust upon a ring-stand and attach a delivery 
tube. Warm the flask gently and collect a jar of the gas by upward 
displacement, as described in appendix. To tell when the flask is 
filled, hold near the mouth a piece of red litmus paper, as in the pre- 
ceding experiment. Keeping the bottle inverted, insert a burning 
taper up into the bottle. Does ammonia burn? Does it support 
combustion ? 

13. Peculiarities of Ammonia. — Ammonia is a colorless 
gas having a strong pungent odor, and if inhaled in con- 
siderable quantities produces strangulation and fills the 
eyes with tears. It is lighter than air, having a density of 
0.59 ; it will not support combustion, nor burn in the air; 
but in oxygen a jet if ignited will continue to burn for 
some time with a yellow flame. It has remarkable affinity 
for chlorine, as will be seen when Ave come to study that 
gas. It also combines readily with hydrochloric acid, 
forming dense white fumes. This will be noticed if two 
bottles, one of each, be opened close together. It is well 
shown also in the following experiment. 

Experiment 44. — Put into a bottle two or three drops of strong 
hydrochloric acid and cover with a glass or paper. Now fill another 
bottle with ammonia gas and invert over the bottle containing the acid. 
Remove the cover separating the two and notice the results. 

14. Solubility in Water. — Ammonia is very soluble in 
water, as the following experiments will show. 

Experiment 45. — Fill the bottle again with ammonia gas as 
before and place it mouth downward into a basin of water. Let it 
stand two or three minutes and notice whether the water rises in the 
bottle. 



NITROGEN AND ITS COMPOUNDS 



77 



15. Ammonia Fountain. — The most striking illustration 
of the solubility of ammonia in water is the "ammonia 
fountain." 

Experiment 46. — Fit to a round-bo ttonied flask or strong bottle, 
of a gallon capacity or more, a rubber cork through which passes a 
long glass tube that will reach half way to the 
bottom of this flask and nearly to the bottom 
of another similar one. Draw out the upper 
end to a jet. Fasten in position or hold over 
the lower bottle or jar as shown in the figure. 
To the water in the lower flask add a few 
drops of some acid and a little litmus solution, 
or a few drops of phenol-phthalein solution. 
Xow fill the upper flask with ammonia as 
in Experiment 44, or by warming gently a 
solution of strong aqua ammonia — the latter 
will be much quicker — and collecting by up- 
ward displacement as before. When well 
filled, quickly insert the cork and long jet- 
tube and support upon the other flask of 
water, as shown in Fig. 19. In a few sec- 
onds the water will begin to rise in the 
tube, owing to the gradual absorption of 
the ammonia, and will soon flow into the 

upper flask. The absorption then will be very rapid, and the water 
will be forced up, forming a beautiful fountain. As it enters the 
upper flask it will change in color, owing to the effect of the ammonia 
upon the litmus or the phenol put into the solution. 

16. Effect of Heat on the Solubility of Ammonia. — At 
0° C. one liter of water will absorb about 1150 liters of 
ammonia. As the temperature rises, the amount absorbed 
rapidly decreases. This is seen in the fact that if a few 
cubic centimeters of aqua ammonia in a flask be warmed 
gently, the gas bubbles out so rapidly that the liquid 
seems to be boiling vigorously when it scarcely feels 
more than warm to the hand. 




Fig. 19. — Ammonia 
Fountain. 



78 



MODERN CHEMISTRY 



17. Effects of Platinum and Charcoal on Ammonia. — 

If a small flask containing strong ammonium hydroxide be 
warmed gently, and a spiral of platinum wire previously 
heated to redness be held in the neck of the flask, the wire 
will continue to glow for a considerable time. Ammonia 
is also absorbed rapidly by charcoal. 

Experiment 47. — To show absorption 
of ammonia by charcoal. 

Fill a large test-tube with ammonia and 
place it inverted over an evaporating dish 
containing a quantity of mercury, as shown 
in the figure. Slip under the tube a piece 
of charcoal. In two or three minutes the 
mercury will begin to rise in the tube to 
fill the space formerly occupied by the 
gas. 




Fig. 20. 



18. Uses of Ammonia. — Immense quantities of ammonia 
are manufactured and used annually. For cleansing pur- 
poses and for softening or " breaking " water it is found 
in almost every household. In a medicinal way it is used 
as a restorative in cases of fainting, and overdoses of chlo- 
roform and other anaesthetics. Considerably diluted it 
is employed in neutralizing the effects of acids upon the 
clothing or upon the hands and face ; in a similar way, 
by inhaling it cautiously, it will counteract the effects of 
chlorine, bromine, sulphur dioxide, and similar irritating 
gases. 

19. As a Refrigerant. — Perhaps the most extensive use 
of ammonia is as a refrigerant in the manufacture of ice. 
The principle underlying this process is as follows : Am- 
monia may be readily liquefied by moderate pressure ; if 
this pressure is suddenly removed, very rapid evaporation 
takes place, producing a low degree of cold. 



NITROGEN AND ITS COMPOUNDS 



79 



Experiment 48. — To show the freezing of water by rapid evapora- 
tion. Put upon a block of wood a few drops of 
water, and upon this a watch crystal. Into the 
crystal pour 1 or 2 cc. of carbon disulphide, and 
blow a current of air by means of a blowpipe 
across the liquid. By the time the disulphide is 
all evaporated the crystal will be frozen tightly 
to the block. Lift the block by taking hold of the crystal. 




Fig. 21. 



20. Manufacture of Ice. 

that ice is manufactured, 
purpose was devised by 



— It is upon the same principle 

The first apparatus for this 

Carre, and is shown in the 




I II 

Fig. 22. — Carre's Apparatus. 



accompanying figure. (I) a is a tank containing strong 
ammonia water, underneath which a fire is placed. This 
causes the ammonia to bubble out of the water very 
rapidly, whereupon it flows over into J, and there liquefies 
by its own pressure, the water surrounding it keeping it 
cool, c is a cylinder of pure water fitting into b. 

21. After a half hour or so, when the ammonia has 
about all been driven out of the solution in a, the posi- 
tion of the two, a and 6, is reversed (II). A partial 



80 



MODERN CHEMISTRY 



r= 




% 


c 




WfrQtti 


A B 






M% 











vacuum forms in a as it cools, the ammonia in b begins 

to evaporate to fill the vacuum, 
and as fast as it flows over 
into a is absorbed by the water 
there. The rapid evaporation 
is thus kept up for a consider-, 
able time, and the cylinder c, 
containing pure water sur- 
rounded by the ammonia 
chamber 5, has its contents 
frozen. 

22. Manufacture of Ice for Commerce. — The first ice 
machines used for the manfacture of ice upon a large scale 
were made upon this principle. At present, however, in- 
stead of creating a vacuum by cooling A, pumps are used 
to remove the vapor from B as fast as it forms. This 
causes, as in the other class of ice machines, a rapid evap- 
oration and a consequent cooling of the adjacent water. 



Fig. 23. — Cross-sectional View of 
Carre's Apparatus. 




Fig. 24. — Modern Ice Plant. 



23. Figure 24 will show the essentials of the improved 
methods of ice manufacture. A is a strong cylindrical 
tank containing liquid ammonia. C is a large rectangular 
vat filled with strong salt water, through which are coiled 
a series of pipes, xx, which connect with A, Through the 



NITROGEN AND ITS COMPOUNDS 81 

top of this vat are let down oblong galvanized iron boxes 
containing the water to be frozen. They are thus sur- 
rounded by the salt water through which the ammonia 
pipes, xx, pass. P is a pump worked by steam, which is 
continually exhausting the pipes and keeping up a rapid 
evaporation in A. The pump, at the same time that it 
exhausts xx, is also condensing the ammonia again in the 
tank M, from which, at intervals, it is allowed to flow 
back again by the pipe y into A. In this way the ammo- 
nia is used over and over without appreciable loss. The 
rapid evaporation lowers the temperature of the salt 
water in below the freezing point of pure water, and 
in from 36 to 60 hours the ice is ready to be drawn from 
the boxes. 

24. Oxides of Nitrogen. — There are five of these com- 
pounds, though not all are of much importance. They 
are : — 

Nitrous Oxide, Laughing Gas, or Nitrogen Monoxide, N 2 

Nitric Oxide, Nitrogen Dioxide, N 2 2 

Nitrous Anhydride, Nitrogen Trioxide, .... N 2 s 
Nitrogen Peroxide, Nitrogen Tetroxide .... N 2 4 
Nitric Anhydride, Nitrogen Pentoxide, .... N 2 5 

The formulae for the second and fourth, for the sake of 
simplicity, are frequently written NO and N0 2 . 

25. Nitrous Oxide. — This is ordinarily called u laugh- 
ing gas." It has been stated already that many of the 
compounds of nitrogen are unstable. So, if we heat am- 
monium nitrate, NH 4 N0 3 , it first melts, then begins to 
boil, and is decomposed to form nitrous oxide and water, 
thus : — 

NH 4 N0 3 + heat = N 2 + 2 H 2 0. 



82 MODERN CHEMISTRY 

Experiment 49. — Put into a test-tube 1 or 2 g. of ammonium 
nitrate, attach a delivery tube, and suspend upon an iron ring-stand. 
Heat moderately and collect two or three small bottles of the gas over 
warm water. Be careful not to heat so strongly as to cause a vigorous 
ebullition, lest some of the impurities always present in the nitrate 
may be carried over and thus vitiate the nitrous oxide. When 
two or three bottles of the gas have been collected, remove the 
cork and notice the odor. Has the gas any color? Test a bottle 
of it with a glowing pine splinter as you did the oxygen. What 
are the results ? Try also a small piece of phosphorus ignited ; how 
does it burn ? 

26. Peculiarities of Nitrous Oxide. — Laughing gas is 
colorless, somewhat heavier than air, having the odor of 
sugar when being heated or slightly burned. It is solu- 
ble to a considerable extent in cold water, will not burn, 
but supports the combustion of most bodies nearly as well 
as oxygen. Upon the human system it acts as an intoxi- 
cant, producing first a sense of hilarity, and after w^ard 
unconsciousness. Because of this fact it is frequently- 
used in a purified form as an anaesthetic in dentistry. It 
is easily liquefied by cold and pressure, and is generally 
used in this form. 

27. Nitric Oxide, N 2 2 . — This gas is almost always one 
of the products formed when a metal is treated with nitric 
acid. 

Experiment 50. — Into a flask put 2 or 3 g. of copper turnings, 
and make connections as for collecting oxygen over water. Add a 
few cubic centimeters of nitric acid, somewhat diluted. What kind 
of fumes first fill the flask? Notice that they disappear, being carried 
over and dissolved in the water. Collect three or four bottles of the 
gas. What can you say of its color and density? Test it to learn 
whether it will burn. Try a blazing pine splinter, also a burning 
candle in the gas ; do they continue to burn ? Try also in a deflagrat- 
ing spoon a well-ignited piece of phosphorus ; what results ? Can you 
explain ? 



NITROGEN AND ITS COMPOUNDS 83 

28. Peculiarities of Nitric Oxide. — As seen above, it is 
a colorless gas, heavier than air, is non-combustible and a 
non-supporter of ordinary combustion. It is noticed, 
however, that substances which burn with great heat, 
such as phosphorus, sodium, and the like, continue to 
burn in nitric oxide with great brilliancy. The reason 
is apparent. Ordinary air is about 20 per cent oxygen ; 
nitric oxide is about 50 per cent oxygen ; such bodies 
therefore as have sufficient heat in burning to decom- 
pose the gas continue to burn more brilliantly, while 
those which kindle at a low temperature have not the 
power to use the large proportion of oxygen present. 

29. Affinity for Oxygen. — The strongest chemical prop- 
erty of the gas is its great affinity for oxygen. This is 
seen whenever it is allowed to escape into the air, brown 
fumes of nitrogen tetroxide being formed. 

Experiment 51. — Into a bottle of nitric oxide inverted over a 
basin of water, pass slowly a current of oxygen. This may be gener- 
ated in a test-tube by using a small amount of potassium chlorate and 
manganese dioxide, or by treating the latter with sulphuric acid. 
Notice how the colorless gas changes ; what else happens ? 

30. One molecule of nitric oxide unites with one of 
oxygen, 2 , as follows : — 

N 2 2 + 2 = N 2 4 . 

If two or three drops of carbon disulphide be put into a 
bottle of nitric oxide and allowed to stand a few minutes, 
or until the disulphide vapor has filled the bottle, on the 
approach of a flame the mixture of gases will burn with a 
brilliant flash, pale violet in color. 

31. Nitrous Anhydride, N 2 3 . — The term, anhydride, 
means icithout tvater, and is applied to certain oxides, 
which, when water is added to them, form acids ; that 



84 MODERN CHEMISTRY 

is, the anhydride is the acid without the water. Such 
oxides were formerly called acids, and carbon dioxide is 
sometimes even now spoken of as carbonic acid ; but all 
true acids contain hydrogen, and theoretically at least 
are formed by adding water to the oxide or anhydride. 
Nitrogen trioxide is thus the anhydride of nitrous acid, and 
is of interest to us only because of this fact. Thus : — 

N 2 3 + H 2 = 2 HN0 2 . 

In this case if the oxide is passed into water it is readily 
absorbed, forming nitrous acid, as shown by the reaction. 

Experiment 52. — Into an evaporating dish put a little starch, 
and with a little water rub it to a thick paste. Transfer this to a 
test-tube into which you have put about 2 cc. of nitric acid. Attach 
a delivery tube and let the end dip into 15 or 20 cc. of water in a 
bottle or flask. Heat the starch in the test-tube for some time, or 
until the fumes are given off readily. What color are they? When 
the gas ceases to come over, test the solution in the bottle with blue 
litmus paper to learn whether it is acid in character. You should 
thus obtain nitrous acid from the brown fumes of nitrogen trioxide 
which were driven off. 

Let us test it to determine. Put one or two cubic centimeters of 
the solution into a test-tube and add a few drops of a solution of 
ferrous sulphate, made by dissolving a crystal of the salt in water. 
Does it turn brown in color? If so, nitrous acid is indicated. 

32. Instability of Nitrous Acid. — Although it is char- 
acteristic of nitrogen compounds to decompose readily, 
nitrous acid is more unstable than most of the others. In 
fact, it breaks up of its own accord at ordinary temperatures. 

Experiment 53. — Put a part of the nitrous acid prepared above 
into a test-tube, and when it has been standing a few minutes, or 
when gently warmed, hold a sheet of white paper behind the tube and 
notice carefully whether brown fumes are being given off. Continue 
to heat gently for a few minutes, or until these do not seem to appear, 
and test with blue litmus paper. Is the solution still acid? If so, test 



NITROGEN AND ITS COMPOUNDS 



85 



a part of it to determine whether it is still nitrous acid. If it is, warm 
it a little longer and test again. If not, test it for nitric acid. This 
is usually done thus : To about 1 cc. of the solution to be tested add 
about as much strong sulphuric acid; shake the two together and cool 
well by holding the tube in a stream of cold water. 

Next prepare a fresh solution of ferrous sulphate, and pour it very 
cautiously upon the solution to be tested so as not to mix them. To 
do this it will be necessary to hold 
the two tubes almost in a horizontal 
position, as show r n in the figure, and 
let the ferrous solution run slowly 
upon the other. Set aside the tube 
in a vertical position and let it 
stand for a few T minutes, w T hen a 
dark brown ring should have 
formed at the junction of the two 
liquids. The test requires consid- 
erable care, but is very satisfactory 

when well done. The test is not distinctive, however, if nitrous acid 
is present, as this also will form a ring ; but the latter ring is usually 
much broader and forms much more quickly. A simpler plan is to 
drop a crystal of ferrous sulphate into the solution to be tested, and 
then pour down the side of the tube upon it a little sulphuric acid. 
A brown ring forms about the ferrous sulphate. 




33. Nitrogen Tetroxide, N 2 4 . — This gas is of little 
importance, yet it is one frequently seen in the action of 
nitric acid upon metals in the presence of air. We noticed 
that when nitric oxide was exposed to the air it quickly 
turned brown. So when a metal is treated with ordinary 
nitric acid, the nitrogen dioxide at first formed quickly 
unites with oxygen from the air and forms the brown 
fumes of nitrogen tetroxide. For experimental purposes 
it may be prepared directly by heating almost any nitrate. 

34. Characteristics of Nitrogen Tetroxide. — It is at ordi- 
nary temperatures a brownish red gas, heavier than air, 
having a very offensive suffocating odor; is non-com- 



86 MODERN CHEMISTRY 

bustible and a non-supporter of ordinary combustion. It 
is soluble in water, as may be seen by inverting a bottle 
of it over water. The brown fumes will disappear and 
the water will rise in the bottle. 

35. Nitrogen Pentoxide, N 2 5 . — The only fact of inter- 
est in connection with this compound is its relation to 
nitric acid, of w^hich it is the anhydride. Hence it is 
often called nitric anhydride. The relation is exhibited 
in the following reaction : — 

N 2 6 + H 2 = 2 HN0 3 . 

36. Nitric Acid, HN0 3 . — When a strong electrical dis- 
charge takes place in the air, as in the case of violent 
thunder storms, small quantities of the nitrogen and oxy- 
gen are caused to unite, forming nitrogen oxides, which 
dissolved in the falling rain form nitric acid. This is 
sometimes in appreciable quantities. Compounds of nitric 
acid, such as sodium and potassium nitrate, are found in 
abundance, especially the former. These salts are now 
known to be produced by the action of certain bacteria 
upon nitrogenous matter, and in some countries the sodium 
nitrate needed is prepared by introducing these bacteria. 

37. Formation of Nitric Acid. — Nitric acid may be ob- 
tained by decomposing any nitrate with sulphuric acid. 

Experiment 54. — Put into a retort 4 or 5 g. of sodium nitrate, 
NaN0 3 , cover with concentrated sulphuric acid, and insert the long 
neck of the retort into a flask surrounded with ice and salt, as in 
Fig. 26. Instead, the retort may be connected with a short Liebig 
condenser, kept cool by a stream of water. Apply a moderate heat 
to the retort ; nitric acid will distil over and condense in the receiver. 
The reaction may be represented in two ways, according to the amount 
of Chile saltpeter used : — 

2 NaNOg + H 2 S0 4 = Na 2 S0 4 + 2 HN0 8 . 

NaN0 3 + H 2 S0 4 = NaHS0 4 + HN0 3 . 



NITROGEN AND ITS COMPOUNDS 



87 




Fig. 26. — Preparation of Nitric Acid. 



38. It will be remembered that we prepared ammonia, 
a volatile compound, by treating a salt of ammonia with 
caustic lime, a compound of similar properties which is 
not volatile. In exactly the same way nitric acid may be 
easily expelled from a liquid by heating, while sulphuric 
acid cannot be. The latter, therefore, simply takes the 
place of the former in combination with the metal, and 
the nitric acid boils out and condenses in the receiver. 
This is shown in the above reactions. 

39. Characteristics of Nitric Acid. — Aqua fortis, as this 
acid is frequently called, is colorless when pure, though, 
owing to impurities present, it is generally slightly yellow- 
ish in color. It is a volatile acid and gives off fumes 
which are very irritating. It colors the skin and finger 
nails yellow, and the color is intensified rather than 
removed by the application of ammonia. Like other 
strong acids, it attacks all organic matter, rapidly destroys 
the fibres of clothing, and the discoloration of the cloth 
cannot be removed by the application of any alkali, as 
is the case with other acids. Though a comparatively 
stable compound, a flask of it exposed to bright sunlight, or 



88 MODERN CHEMISTRY 

heated, soon becomes filled with a brownish gas, nitrogen 
peroxide, N 2 4 , and oxygen. This is seen in the follow- 
ing reaction : — 

2 HN0 3 + heat = O + H 2 + N 2 4 . 

On account of this property of giving up a part of its 
oxygen with moderate ease it is frequently used as an 
oxidizing agent. This is seen in the following experi- 
ments : — 

Experiment 55. — Warm slightly a little turpentine in an evap- 
orating dish and pour upon it some strong or fuming nitric acid. 
Usually only a copious evolution of fumes is the result, but sometimes 
the oxidation is so rapid that the whole mass bursts into a flame. 

Experiment 56. — Heat in an iron spoon a few small pieces of 
charcoal ; when red hot drop quickly into a beaker containing some 
strong nitric acid. Notice that the charcoal continues to glow for 
some little time, owing to the oxygen obtained from the acid ; notice 
also that brown fumes fill the beaker. Upon a small quantity of 
warm strong nitric acid in an evaporating dish, drop a very small 
piece of phosphorus. Notice that it is instantly set on fire, and small 
particles are thrown out in all directions. 

Experiment 57. — To a little tin-foil in a test-tube add some strong 
nitric acid and heat. Notice that the metal is not dissolved, but con- 
verted into a white solid, which is really an oxide of tin in combina- 
tion with water, Sn0 2 , H 2 0. By heating, the water is evaporated 
and the oxide remains. 

40. Uses of Nitric Acid. — Nitric acid finds a great 
many uses in the laboratory, frequently as an oxidizing 
agent, as will be seen from time to time. It is used con- 
siderably in the manufacture of sulphuric acid, which will 
be described later, and in making nitro-glycerine and other 
explosives. 

41. Aqua Regia. — This is a mixture of nitric and hydro- 
chloric acids in the proportion of one of the former to 
three of the latter, and is so named because it will dis- 



NITROGEN AND ITS COMPOUNDS 89 

solve gold, the "king of the metals." It is the strongest 
solvent known, and attacks several metals which are 
unaffected by single acids. 

42. Nitro-glycerine and Dynamite. — Nitro-glycerine is 
prepared by treating glycerine with a mixture of fuming 
nitric and sulphuric acids. It is in the form of a liquid, 
and hence not convenient for uses under all circumstances. 
Dynamite differs from nitro-glycerine in that it contains 
about 25 per cent of siliceous or infusorial earth. It is, 
therefore, more convenient and less liable to explode by 
accident. Guncotton is a similar compound which is pre- 
pared by treating cotton wool with nitric and sulphuric 
acids. It is, therefore, not very different from nitro-glycer- 
ine in composition. It has the advantage of being perfectly 
safe when wet, and is, therefore, kept damp when carried 
on board men-of-war. In this condition it is exploded 
by igniting with a small charge of fulminating mercury. 
Its combustion is five hundred times as rapid as that of 
the best gunpowder. The heavy charges now used for tor- 
pedoes give an impact that no man-of-war can withstand. 
All of these explosives, as well as gunpowder, are valuable 
because of the great instability of the nitrates present or 
formed in the preparation of them. 

Argon : A = 40 ? 

43. Its Discovery. — For some time previous to the dis- 
covery of argon, in 1894, it had been observed that nitro- 
gen obtained from the atmosphere was heavier than that 
from its compounds. In that year Lord Rayleigh and 
Professor Ramsay observed that, by passing atmospheric 
nitrogen over red-hot magnesium, a small residue was 
obtained which could not be made to enter into combina- 
tion. This residue was the n^y gas now called Argon, 



90 MODERN CHEMISTRY 

Its name comes from the Greek word argon^ which means 
idle or inactive. 

44. Characteristics of Argon. — This element is an odor- 
less, colorless gas, somewhat heavier than air, constituting 
about eight-tenths of one per cent of the atmosphere. As 
far as is known it is a perfectly inert substance, hitherto 
resisting all attempts to make it enter into combination. 
No compounds of the gas being known, it is impos- 
sible to assign it a positive atomic weight, but it is 
believed to be about forty. 

SUMMARY OF CHAPTER 

Origin of the term nitrogen. 

Abundance of the element and of certain compounds. 

Easiest method of preparing nitrogen. 

What is the purpose of the phosphorus ? 
The source of the nitrogen ? 

Would a candle do as well as phosphorus? Why? 
Two other ways of preparing nitrogen. 

Chemical action in each case. 
Characteristics of nitrogen. 

Compare with oxygen and hydrogen — How similar — How 

different. 
How test each ? 
Compounds of nitrogen. 

Ammonia — How formed in nature. 

Old method of preparing " hartshorn." 
Present source of ammonia. 
Wherein are these three methods similar ? 
Characteristics of ammonia. 

Experiments to illustrate these. 
Uses. 
Experiments to illustrate the most impor- 
tant. 
Carre's ice machine. 
Present ice machines, 



THE ATMOSPHERE 91 

Oxides of nitrogen. 

Names and formulae. 

Most important. Why? 

Method of preparing this one. 

Use — Physiological effects. 
Acids of nitrogen. 

Names and formulae. 

Anhydride of each — Meaning of. 

How distinguish each by test. 

Preparation of nitric acid. 

Characteristics and uses of. 

Aqua regia. 

Explosives — Explanation of their explosive character. 



CHAPTER VIII 

THE ATMOSPHERE 

1. What it is. — We are living at the bottom of an 
ocean as wonderful as the watery one that washes the 
shores of our continent. The atmosphere covers the 
entire earth to a depth variously estimated at from fifty 
to two hundred miles. Some of the recent investigators 
believe that, in an extremely attenuated form, the air 
extends through space, even reaching and commingling 
with the atmospheres of other planets. Centuries ago 
the air was regarded as one of the elements, just as water 
was, and the other gases, as discovered, were all called 
air ; for example, hydrogen was known as inflammable air, 
carbon dioxide as fixed air, etc. So the perfumes that 
were exhaled from various flowers were regarded as air, 
slightly changed in some unknown way. 

2. Constituents of the Air. — We know now that the 
air is not an element, but a mixture of several substances. 
Three of these, nitrogen, oxygen, and argon, are con- 



92 MODERN CHEMISTRY 

stant, but the watery vapor and carbon dioxide vary from 
time to time. Many efforts have been made to learn 
whether the air is a compound of oxygen and nitrogen, 
mixed with the other constituents named. Analyses have 
been made in ail parts of the world, thousands of feet 
above the earth, in the crowded cities, on the North and 
South American prairies ; but though the proportion of 
the gases, 79 of nitrogen to 21 of oxygen, by volume, is 
found in all cases approximately the same, yet the varia- 
tion is too great to permit one to believe that they are 
united to form a compound. The argon constitutes about 
1 per cent of what has usually been taken as nitrogen, or 
about 0.8 per cent of the air. The carbon dioxide varies 
somewhat, but seldom amounts to more than three or four 
parts in 10,000, except in poorly ventilated rooms. The 
aqueous vapor varies greatly. When the air contains all 
it is able to hold, it is said to be saturated, or to contain 
100 per cent. Ordinarily, however, the humidity is not 
above 60 to 70 per cent. The amount may be estimated 
by passing a certain volume of air through a tube filled 
with calcium chloride and noting the increase of weight. 

3. Diffusion of Gases. — We find that the air contains 
five gases, of densities ranging all the way from eight to 
forty times that of hydrogen. Were it not for the law 
of diffusion, we should find the argon, perhaps, nearest 
the ground The next above this, forming a layer twelve 
feet or more deep, would be the carbon dioxide ; then 
the oxygen, nitrogen, and water vapor, in the order 
named. Such conditions would be fatal to all animal 
life. As it is, however, owing to the constant circula- 
tion of the atmosphere and the rapid diffusion of gases, 
no more carbon dioxide is found close to the surface of 
the earth than hundreds of feet above. Two or three 



THE ATMOSPHERE 



93 



exceptions to this ought to be noted, among them the 
deadly Upas Valley, where the carbon dioxide is exhaled 
from volcanic sources more rapidly than diffusion can 
carry it away.* 

4. Boyle's Law. — Many years ago Boyle discovered 
and formulated the law, which now bears his name, 
that the volume of a gas, the temperature 
remaining constant, varies inversely as the 
pressure. In other words, if we double 
the pressure, the volume decreases by 
half ; or if we lessen the pressure by half, 
the volume becomes twice as great. In the 
accompanying figure we have 10 cc. of the 
gas, a, under the pressure of the atmos- 
phere, simply confined by the mercury in 
the bottom of the bent tube ; if now we pour in more 
mercury at the open end, the volume of a will constantly 
decrease, and when we have added as much as corre- 
sponds to the pressure of an additional atmosphere, the 
volume will have decreased to 5 cc. 

5. Standard Pressure. — Atmospheric pressure is meas- 
ured by the barometer, which at sea level stands about 
30 in. high. In chemical calculations, however, we use 
the metric system, and the equivalent of 30 in. is 760 mm. 
Hence when we say that a gas is under standard pressure, 
we mean 760 mm. 




Fig. 27. 



* This valley is located in the island of Java, is about a half mile in 
circumference and thirty-five feet deep, surrounded at no great distance 
by hills. The bottom is comparatively smooth and is devoid of vegeta- 
tion. Loudon, in describing his visit there, says that "skeletons of human 
beings, tigers, pigs, deer, peacocks, and all sorts of birds " are to be seen 
everywhere, bleached by the exposure till they are as white as ivory. A 
fowl thrown in died in one and a half minutes. 



94 MODERN CHEMISTRY 

Problem. — 500 cc. of oxygen under standard pressure would be 
how much under 750 mm.? As the pressure has decreased, the 
volume would have increased. We would solve then by the fol- 
lowing proportion : — 

V: V'\\P'\P\ 

or 500 cc. : x : : 750 : 760. 



If desired, this problem may be solved without using a proportion. 
As the pressure has decreased, we know that the volume will be corre- 
spondingly increased ; that is, 

V will equal ff§ of V; 

or V = ffg x 500. 

V = ? 

2. What volume would 300 cc. of hydrogen, at 750 mm. pressure, 
occupy at 780 mm. pressure ? 

3. 25 liters of air at 380 mm. pressure, would be how many at 5 
atmospheres' pressure ? 

6. Law of Charles. — Just as heat causes solids and 
liquids to expand, so it affects gases. In the case of 
the latter the rate of expansion is practically constant, and 
is in the Law of Charles stated thus : The pressure re- 
maining constant, all gases expand or contract uniformly 
under the same increase or decrease of temperature. This 
has been studied carefully, and it has been proven that 
for an increase or decrease of 1° C, a volume of gas ex- 
pands or contracts ^g of the volume it occupies at 0° C. 
To illustrate, suppose we have in a vessel 273 cc. of 
oxygen at 0° C. If by any means the temperature is 
raised to 10° C, the volume would increase ^V ^ of 273 cc. 
or 10 cc, and would occupy 283 cc. It obviously follows 
from this that were the law to hold true, and were the gas 
reduced to a temperature of 273° below zero, it would 
disappear entirely. However, all gases thus far tried 



THE ATMOSPHERE 



95 



become liquids before reaching this low temperature, so 
that the law no longer applies. 

7. Absolute Zero. — From the fact that a gas would 
disappear entirely at 273° below zero, according to the 
Law of Charles, — 273° has been called absolute zero, the 
point at which the molecules of a body would have no 
vis viva, or absolutely no heat energy. This point has 
never yet been reached, though recent investigators have 
approached within a few degrees of it. It is necessary to 
have a clear understanding of what is meant by the abso- 
lute zero, as it is used in making all calculations for 
correction of the volume of gases for temperature. 

8. The Absolute Thermometer. — In Fig. 28 we have 
the Fahrenheit, Centigrade, and Absolute thermometers 
represented in F, C, and A, 

respectively. It must be re- 
membered that the last is not a 
thermometer really in existence, 
but serves merely for illustra- 
tion. The boiling points on the 
three are marked 212°, 100°, and 
373°; the freezing points 32°, 
0°, and 273°. The absolute zero 
therefore would be the same as 
— 273° on the Centigrade, as 
the degrees on these two are of 
the same size. Let us apply 
this in a problem. 



ZJZ' 



3Z l 
0° 



100- 



373 



Z73 



# 273 # "# 



-Boiling £oint 



—Freezing PL 



Fig. 28. 



Problem. — 500 cc. of oxygen at 0° C. would occupy what volume 
at 25° C. ? 

Expressing Charles's Law in the form of a proportion, we would 
have 

V: V'::t:t f , 



96 MODERN CHEMISTRY 

in which V and t represent the volume and temperature of the gas at 
the beginning of the experiment, and V and t' at the end; and it 
must be remembered that t and t' always mean absolute temperatures. 
Applying this to the problem, we see that t = 0° C, or 273° A., and 
? = 273 + 25, or 298° A. Substituting, we have : — 

500 : V : : 273 : 298. 
v , = 298 x 500 
273 

Problem 2. — What volume would 250 cc. of gas at 20° C. occupy 
at -10° C.? 

Here, t = 20° C. = 273 + 20 = 293° A. 

t > = _ io° C. = 273 - 10 = 263° A. 
F=250cc. 
Substituting, 250 : V : : 293 : 263. 

y, = 2o0 x 263 
293 

Problem 3. — If 400 cc. of hydrogen is heated from — 15° C. to 
30° C, what volume would the gas then occupy? 

Problem 4. — What would be the result in problem 2, if at the 
same time the barometer fell from 760 mm. to 740 ? 

This may be solved by first finding the value of V, as shown above 
at the temperature t f , and substituting this in the proportion for 
determining V under P' pressure. Suppose, in problem 2 above, 
F' = 225+, then solving for pressure, we would have 

V" = Ifg of 225 + ; 
or V : V" ::P ff :P f ] 

or 225 : V" : : 740 : 760. 

y tt = 225 x 760 
740 

Or the problem may be solved by using a compound proportion : — 

( 97S 4- 90 • 273 — 10 
250 :x : it ~ /" ; 293 x 740* = 263 x 760 x 250. 

Problem 5. — 500 cc. of gas under 4 atmospheres and at — 25° C. 
would have what volume at 760 mm. and at 20° C. ? 

Let the teacher furnish a number of similar problems for practice. 



THE ATMOSPHERE 



97 



9. Weight of Air. — The weight of a liter of air may 
easily be found by the following experiment : — 

Experiment 58. — M in the figure is a flask 
of about 500 cc. capacity. Fit to it a cork with 
a glass tube somewhat drawn out, as shown. 
Put into the flask about 50 cc. of water and 
boil for several minutes, so as to expel all the 
air. Immediately remove the cork and insert 
another, not perforated. When the flask has 
cooled to the temperature of the room, weigh 
the whole. Suppose this to be m. Remove the 
cork, thus allowing the air to enter, and again 
weigh flask and cork. Suppose this to be n. 
The gain in weight, 

n — m = wt. of air in flask. 




Fig. 29. 



To determine the volume of the air contained, take a graduated flask, 
or cylinder, and fill the flask M with water. Suppose this to be r cc. 
Then 

r cc. of air weighs n — m grams, 

from which the weight of 1000 cc. = 1 liter may be determined. 

10. Liquefaction of Air. — The air is so well known that 
it is not necessary to say anything regarding its proper- 
ties. At the present time, however, considerable atten- 
tion is being given to it in the liquid form. A large 
number of experiments with it have been made by Dewar, 
Pictet, Linde, Tripler, and others, with a view to ascer- 
taining its properties and practical value. It is said that 
the first ounce of liquid air ever produced cost about 
13000 and the next pint about $80 ; with improved 
methods, however, it may now be prepared for a few cents 
per gallon. 

11. Dewar's Bulbs. — Dewar has invented a double- 
walled glass globe in which liquid air may be kept for 
a number of hours with little loss ; here in this country 





98 MODERN CHEMISTRY 

it is often shipped several hundred miles in large double- 
walled tin cans, heavily lined with felt, but at the 
expense of 20 to 40 per cent of the liquid. The Dewar 
bulbs vary somewhat in construction, but the general 

plan is the same in all. Into 
the space between the inner 
and outer walls of the globe, 
a drop or two of mercury is 
introduced; the air-pump is 
then attached, and a vacuum 
of very high degree obtained. 
As the air is pumped out, the 
mercury vaporizes and fills the 
space. When liquid air is introduced into the inner globe, 
the mercurial vapor is condensed upon the outer surface 
of the inside flask, and forms a perfect mirror. Thus Ave 
have not only a vacuous chamber, but also a mirror to 
prevent the access of heat rays to the liquid air, and the 
insulation is well nigh perfect. A modified form of 
this Dewar bulb, holding about two gallons, is now used 
for shipping liquid air. The insulation is so perfect 
that the liquid may be kept two weeks with little loss. 
It is obvious that the ordinary closed tank is unsuit- 
able on account of the high pressure which would soon 
obtain. 

12. Linde's Apparatus. — The plan used for liquefying 
air may be understood from the accompanying figure, 
which represents the apparatus used by Linde. P is a 
pump which, when the piston is raised, opens a valve at 
Gr and allows the air from D to enter; as the piston 
descends, the valve Gr closes and IT opens. The air is 
thus forced up through the coils in the tank J, kept cold 
by running water, and passes on through B. At O the 



THE ATMOSPHERE 



99 



pipe B enters within a larger one, and continues thus 
until at the point E it again emerges. The ingoing cur- 
rent of air flows through the inner pipe under pressure 
and issues from a small aperture at R into a chamber, T, 
under low pressure. As expansion is a cooling process, 
the air is thus reduced in temperature ; at the next stroke 




Used by Courtesy of the Scientific American. 

Fig. 31.— Linde's Apparatus for liquefying Air. 

of the piston the vacuum formed in the pump again opens 
the throttle valve at (7, and the cooled air in T flows back 
through the outer pipe, back through D. As this opera- 
tion is constantly repeated, the outgoing current being 
cooled by its expansion into T, continually lowers the 
temperature of the ingoing current, until finally liquid 
air will trickle down into the chamber T, and may be 
drawn off at Fmuch the same as water from a reservoir. 



L.ofC. 



100 MODERN CHEMISTRY 

13. Effects of Liquid Air upon Certain Substances. — It 

is found that such articles as rubber, beefsteak, eggs, etc., 
immersed in liquid air, become exceedingly brittle ; while 
an ordinary tin cup dipped into the liquid and dropped 
upon the floor breaks into fragments like glass. All these 
effects are due to the intense degree of cold of the liquid 
air, and not to any chemical action. 

14. As the boiling point of nitrogen is lower than that 
of oxygen, the former boils out the more rapidly, and 
in a short time a vessel of liquid air, freely exposed, will 
contain almost pure liquid oxygen. If into this a red- 
hot iron rod be thrust, it will burn vigorously, notwith- 
standing the fact that the temperature of the surrounding 
liquid is nearly 1700° C. below the melting point of iron. 
It should be said, however, that the two are probably 
not in contact, but that a layer of gaseous oxygen next to 
the iron rod supports the combustion. Felt, saturated 
with liquid air, burns explosively, and if confined in metal 
tubes, bursts them with violence. 

15. Practical Uses of Liquid Air. — Numerous applica- 
tions for liquid air have been suggested, but as yet these 
are in the experimental stage. Among them may be 
named the following : (1) as a substitute for compressed 
air ; (2) as a refrigerant ; (3) in blasting ; (4) in surgery 
for removing diseased tissues without the use of the 
knife ; (5) as a smoke consumer, and for burning garbage 

in cities. 

SUMMARY OF CHAPTER 

Composition of the atmosphere. 

Old ideas of the air. 

Present ideas. 

Explanation of uniformity of composition. - 
Boyle's Law — Statement of. 

Meaning of term standard pressure. 



THE HALOGENS 101 

Charles' Law — statement of. 

Meaning of term absolute zero. 

Problems. 
Density of air. 

Methods of finding weight of one liter. 
Liquefaction of air. 

Present method. 

Dewar bulbs. 

Effects of liquid air. 

Suggested uses. 



CHAPTER IX 

THE HALOGENS 

1. Members of the Group. — The term halogen is from 
two Greek words, meaning salt producer, and is given to 
this group of elements because with the metals they form 
a large number of salts. The group includes fluorine, 
chlorine, bromine, and iodine. The first two are gases, the 
third a liquid, and the fourth a solid. They possess prop- 
erties very similar to each other, differing in degree 
rather than otherwise. It will be found that as the atomic 
weights increase, the chemical activity decreases. 

Fluorine : F = 19 

2. Characteristics. — Fluorine is an element which had 
not been prepared until a few years ago. It is a greenish- 
colored gas, of a very irritating odor, and readily attacks 
almost all substances. By extreme cold and pressure it 
has been liquefied, and when in that condition loses much 
of its chemical activity. It is of little practical value, and 
is considered only because of one or two compounds which 
it forms. 



102 MODERN CHEMISTRY 

3. Compounds of Fluorine. — There is only one com- 
pound of this element in which we are specially interested, 
and that is hydrofluoric acid, HF. It is prepared by 
treating fluor spar, calcium fluoride, with strong sul- 
phuric acid, the reaction being — 

CaF 2 + H 2 S0 4 = CaS0 4 + 2 HF. 

Hydrofluoric acid is a very irritating, colorless gas, which 
readily dissolves or corrodes glass, and hence is sometimes 
used in. glass etching. 

Experiment 59. — Warm a sheet of glass 3 or 4 in. square by 
holding it at some height above the burner flame, and drop upon it 
a few shavings of parafiine. Move the glass about so as to distribute 
the melted wax evenly, and allow it to cool. Now with a sharp pen- 
cil or stylus draw any desired figure in the wax, being sure to cut 
through to expose the glass. Lay this face down over a lead saucer,* 
into which you have put about 2 g. of calcium fluoride and as much 
strong sulphuric acid. Support upon a ring-stand and warm for a 
minute very gently, so as not to melt the wax. In a few minutes the 
etching should be completed. This can be determined by testing 
with the point of a knife blade, when the glass will feel rough where 
the figure was drawn in the wax. When the experiment is finished, 
the parafiine may be removed with a dull knife or by immersing in 
warm water. 

Chlorine : CI = 35.5 

4. History. — Chlorine, the most important element of 
the halogen group, was first prepared by Scheele in 1774, 
in treating black oxide of manganese — the same com- 
pound we have used in preparing oxygen — with hydro- 
chloric acid. He did not know, however, that he had 
discovered a new element, but supposed it to be a com- 

* Instead of the lead saucer a small evaporating dish may be used. If 
so, notice whether it also is attacked on the inside by the hydrofluoric 
acid. 



THE HALOGENS 103 

pound of oxygen and hydrochloric acid, and called it 
dephlogisticated marine acid air. Hydrochloric acid was 
then called marine acid. Later, when chlorine was found 
to be an element, it was given its present name from 
the Greek word chloros, meaning green. 

5. How found. — Because of its great chemical affinity, 
chlorine, like fluorine, is never found uncombined. Its 
most widely distributed compound is common salt, NaCl, 
which is found in extensive deposits in nearly all parts of 
the United States, and constitutes a large per cent of the 
solids held in solution in the ocean. 

6. How to prepare Chlorine. — For laboratory purposes 
the simplest way of preparing chlorine is that used by its 
discoverer, by treating manganese dioxide with hydro- 
chloric acid and warming gently. 

Experiment 60. — Into a good-sized test-tube or generating flask 
put 1 or 2 g. of manganese dioxide and about 2 cc. of hydrochloric 
acid. Attach a delivery tube and warm gently. Collect two or three 
bottles of chlorine by downward displacement, as described on page 
362. and preserve for future experiments in studying its properties. 

7. The reaction that takes place in preparing chlorine 
as above may be indicated thus : — 

Mn0 2 + 4 HC1 = Cl 2 + MnCl 2 + 2 H 2 0. 

From this we see that only half the chlorine, in the hydro- 
chloric acid used, is obtained free, the other half having 
united with the manganese to form manganese chloride, a 
compound which has no application in the arts. An 
immense quantity of chlorine is used every year in the 
manufacture of bleaching powder, and cost of production 
is a very important consideration. The method described 
above is, therefore, not strictly followed commercially, 
but is so modified that the manganese chloride is converted 



104 MODERN CHEMISTRY 

into the dioxide again. This is much cheaper, and is 
known as the Weldon process. 

8. The Weldon Process. — In the preparation of chlorine 
for manufacturing processes, pyrolusite, a natural ore of 
manganese and an impure form of the dioxide, is treated 
with hydrochloric acid in large stone tanks. When the 
chlorine is no longer given off, any excess of acid in the 
residual liquor is neutralized with common limestone, 
finely powdered. The reaction may be represented thus : — 

MnCl 2 + H 2 + 2 HC1 + CaC0 3 

= MnCl 2 + CaCl 2 + C0 2 + 2 H 2 0. 

(Residual liquor and excess\ , /'limestone') = /mixture manganese and cal-\ 
of acid / \ cium chloride in water. / 

Therefore, we now have a mixture of manganese chloride 
and calcium chloride in solution. Next, lime water, pre- 
pared by treating ordinary lime with water, is added. 

CaO (lime)+ H 2 = Ca(OH) 2 (lime water). 

This precipitates the manganese in the form of the hydrox- 
ide, Mn(OH) 2 , thus : — 

CaCl 2 } + C < 0H >2 = Mn(OH) 2 + 2 CaCl 2 . 

Now by heating this and at the same time passing a cur- 
rent of air through the solution, the manganese hydroxide, 
Mn(OH) 2 , is converted into the dioxide, thus : — 

Mn(OH) 2 + O (air)= Mn0 2 + H 2 0. 

The calcium chloride, being very soluble, remains in solu- 
tion. The mixture is now allowed to flow into settling 
basins, where the dioxide is slowly deposited as a dark- 



THE HALOGENS 105 

colored ooze, known as Weldotis mud. This is now ready 
to be passed again into the stills for a second treatment 
with hydrochloric acid.* 

9. The Chemical Changes in the Above Method. — By 
studying the reaction 

Mn0 2 + 4 HC1 = MnCl 2 + 2 H 2 + Cl 2 , 

we see that the oxygen in the manganese dioxide has been 
set free from the manganese and has united with the hy- 
drogen in the acid ; or, as we sometimes say, the chlorine 
has been set free by the oxidation of the hydrogen with 
which it was combined. In like manner other substances, 
besides manganese dioxide, may be used with hydrochloric 
acid in preparing chlorine. In every instance the princi- 
ple is the same : the oxygen is first set free, and, combining 
with the hydrogen in the acid, liberates the chlorine. Let 
us prove this. 

Experiment 61. — Treat a few crystals of potassium chlorate, 
KC10 3 , a substance from which we obtained oxygen, with a little 
hydrochloric acid. Warm very gently if necessary to start the action, 
and then remove the test-tube from the flame. ^Notice the rapid evo- 
lution of gas. With the chlorine thus obtained we have also an oxide 
of chlorine, C10 2 , as seen in the reaction 

4 KC10 3 + 12 HC1 = 4 KC1 + 9 CI + 3 C10 2 + 6 H 2 0. 

Add to this a few cubic centimeters of water, which will give a 
yellowish solution known as euclilorine or chlorine water. Preserve it 
in a dark-colored, tightly stoppered bottle. 

* It perhaps ought to be stated that a small excess of lime water usu- 
ally remains mixed with the precipitated manganese hydroxide. When 
j the current of air is passed through the solution, this lime water, Ca(OH) 2 , 
is also oxidized ; that is, converted into lime, CaO. Weldon's mud, 
therefore, contains besides the manganese dioxide, a small amount of 
lime. 

\ 



106 MODERN CHEMISTRY 






10. It will be remembered that we prepared oxygen 
also by using potassium dichromate. If now we treat this 
compound with hydrochloric acid, chlorine is obtained as 
in the other instances. 

11. Practical Application of this Principle. — In all the 
above methods the chlorine is set free by bringing into 
contact with hydrochloric acid some highly oxygenized 
substance which will give up a part of its oxygen to unite 
with the hydrogen of the acid. Hence was conceived the 
idea of using atmospheric oxygen as the most economical 
source of supply. 

12. Deacon's Process. — This idea is applied in Deacon's 
process. Theoretically the reaction that takes place ac- 
cording to this method is as follows : — 

2 HC1 + O = H 2 + Cl 2 . 

In reality, however, the process is not so simple. In the 
preparation of oxygen from potassium chlorate and man- 
ganese dioxide, we have seen that the latter compound 
remains unchanged. . It acts, as was said, by catalysis, in 
a manner not thoroughly understood, causing the potas- 
sium chlorate to yield up its oxygen at a temperature 
much lower than would otherwise affect it. 

13. The Catalytic Agent. — In Deacon's process for the 
preparation of chlorine some catalytic agent is necessary, 
because a mixture of oxygen and gaseous hydrochloric 
acid, when heated, is only slightly decomposed. As a 
catalytic agent some such compound of copper as the 
sulphate or the chloride is used. Clay balls or bits of 
brick are saturated with the copper solution and placed in 
an iron pipe called the decomposer. Through this the 
mixed gases, air and hydrochloric acid, previously heated 
to about 500° C, are made to pass. The acid is oxidized 



THE HALOGENS 107 

and the chlorine set free. The chemical action of the 
cuprous chloride, Cu 2 Cl 2 , is not thoroughly understood ; 
but it is believed that two or three reactions take place, 
in the course of which cupric chloride, CuCl 2 , is formed, 
which at the temperature present is unstable and gives up 
a part of its chlorine, leaving cuprous chloride again. 

14. Another Method of preparing Chlorine. — Another 
method is frequently employed in the laboratory instead 
of the first one given. 

Experiment 62. — Into a test-tube put a small quantity of common 
salt, NaCl, mixed with a little manganese dioxide, and about a cubic 
centimeter of sulphuric acid. Warm gently. Is there any evidence 
that chlorine is being generated ? 

15. Comparison of the Two Methods. — We shall find 
that when common salt is heated with sulphuric acid they 
react with each other, forming hydrochloric acid. That is 
what we have done in this case. We see by comparing 
the two reactions, 

Mn0 2 + 4 HC1 = MnCl 2 + 2 H 2 + Cl 2 
and ^ 

Mn0 2 + 2N^Cl + 2H 2 S0 4 =MnS0 4 +Na 2 S0 4 +2H 2 0+Cl 2 , 

that in the first case we treated the dioxide directly w r ith 
hydrochloric acid, but in the second indirectly by the use 
of two substances, which in reacting prepare the hydro- 
chloric acid needed. It will be seen, however, in the sec- 
ond instance that all the chlorine is set free, while in the 
first only one-half. 

16. It is probable that in the second case the reaction 
is a little more complicated, perhaps as follows : — 

First, a part of the sulphuric acid reacts with the com- 
mon salt, forming Irydrochloric acid, thus : — 

2 NaCl + H 2 S0 4 = 2 HC1 + Na a S0 4 . 



108 MODERN CHEMISTRY 

Then another part reacts with the manganese dioxide 
also present, setting free oxygen, as we have seen before, 
thus : — 

Mn0 2 + H 2 S0 4 = MnS0 4 + H 2 + O. 

Then this nascent oxygen immediately attacks the hydro- 
chloric acid present, oxidizing it and liberating the chlorine, 
thus : — 

2 HC1 + O = H 2 + Cl 2 . 

Putting these three reactions together, we would have 

2NaCl + 2H 2 S0 4 + Mn0 2 =MnS0 4 + Na 2 S0 4 + 2H 2 + Cl 2 . 

17. Experiments with Chlorine. — With the chlorine 
prepared make the following experiments in study of its 
properties : — 

Experiment 63. — Note the color of the gas; the odor. Put a 
burning match into a bottle of chlorine ; try also a burning candle. 
State the results. Does the gas burn? Does it support combustion? 

Experiment 64. — To show its chemism for certain metals. Sift 
into a bottle of chlorine, by means of a fine wire-gauze spoon, some 
powdered metallic antimony; try in the same way metallic arsenic. 
Describe the results. 

Experiment 65. — To show the chemism of chlorine for hydrogen. 
In a room partially darkened, fill a strong bottle with chlorine and 
hydrogen, mixed. Wrap a towel about it, ignite a piece of magnesium 
ribbon, and bring it toward the mouth of the bottle. A violent explo- 
sion is the result. Bright sunlight has the same effect. Try also the 
following experiment to show the same fact. 

Experiment 66. — Attach a jet to a hydrogen generator, H, and 
when it has been in action long enough to expel all the air, ignite it, 
and insert into a jar of chlorine, C, as shown in Eigure 32. Does it 
continue to burn ? How does the flame change in appearance ? What 
becomes of the green gas? After a few moments add about 1 or 2 cc. 
of water to the gas, and shake well. Drop into the solution a piece of 
blue litmus paper ; what is indicated ? It is best to dry the hydrogen 
by passing through a drying tube, Z), filled with calcium chloride. 



THE HALOGENS 



109 




Fig. 32. 



Experiment 67. — To show affinity of chlorine for hydrogen in 
compounds of the latter. Into a jar of chlorine thrust a narrow slip 
of blotting paper which has been 
moistened in moderately warm tur- 
pentine. State the results. Turpen- 
tine consists of carbon and hydrogen, 
C 10 H 16 . What has the chlorine really 
done ? 

Experiment 68. — Practical appli- 
cation of the preceding experiment. 
Into a jar of chlorine pour a few cubic 
centimeters of any solution containing 
organic colors, as litmus, logwood, or 
carmine. Shake it up and notice the 
effects. 

Experiment 69. — With the same 
purpose as in Experiment 68. In 
another jar of chlorine, suspend a 
piece of blue or pink calico mois- 
tened with water. Try another simi- 
lar piece without moistening it. Are the results different? 

Experiment 70. — To show affinity of chlorine for ammonia. At- 
tach to a small flask, into which you have put 25 or 30 cc. of strong 
aqua ammonia, a delivery tube with jet attached. Warm the flask 
gently as in preparing ammonia for the " fountain," Experiment 46, 
and insert the tube into a bottle well filled with chlorine. What 
happens ? What becomes of the chlorine ? 

18. Characteristics of Chlorine. — Chlorine is a greenish 
yellow gas, with a very irritating odor, producing tem- 
porarily a catarrhal affection of the nasal passages. It is 
somewhat soluble in water, forming a solution yellowish 
in color, with the characteristic odor of chlorine. This 
solution is, however, unstable, as the chlorine gradually 
combines with the hydrogen of the water to form hydro- 
chloric acid, while the oxygen is set free. Chlorine is 
about two and a half times as heavy as air and does not 
support ordinary combustion. It will be found, however, 



110 MODERN CHEMISTRY 

that sodium and phosphorus, when well ignited, burn vig- 
orously in an atmosphere of chlorine. 

Experiment 71. — Put a small piece of sodium, heated in a defla- 
grating spoon until it takes fire, into a bottle of chlorine. State 
results. Notice the white deposit of common salt that forms. In 
the same way try a piece of phosphorus, without first igniting it. 
State results. 

19. Chemical Affinity of Chlorine. — From our experi- 
ments in oxygen we learned that considerable heat was 
necessary to effect its rapid union with any other element. 
The iron wire, the sulphur, and the phosphorus, all had to 
be raised to the kindling point. In the case of chlorine 
we find that union often takes place at ordinary tem- 
peratures, showing its chemism to be far greater. Thus 
arsenic and antimony sprinkled into the gas took fire 
spontaneously, as did also the phosphorus and the tur- 
pentine. In the latter case the chemical action is due 
to the affinity between the hydrogen in the turpentine 
and the chlorine ; the same remarkable affinity of these 
gases for each other was also seen in exploding the mix- 
ture of the two by means of light, and in the hydrogen 
jet which continued to burn in the chlorine. 

20. Chlorine as a Solid. — If a saturated solution of 
chlorine water be surrounded by a mixture of ice and 
salt, in a few minutes yellowish crystals of chlorine 
hydroxide, represented by the formula CI, 5 H 2 0, are 
formed throughout the liquid. Chlorine may be lique- 
fied at —34° C. under ordinary atmospheric pressure, 
or at 0° with a pressure of six atmospheres. In this 
condition it is of a bright yellow color. It has also been 
solidified by reducing to 102° below zero, and in this form I 
closely resembles the liquid in color. 



THE HALOGENS 



111 



21. Uses of Chlorine. — Chlorine is used to a consider- 
able extent in the extraction of gold from its ores, because 
it is a good solvent of that metal. A large amount of 
that now used for this purpose is put up at the factories 
in the liquid form in steel cylinders lined with lead, and 
then shipped wherever desired. 

22. As a Bleaching Agent. — Chlorine is a powerful 
bleaching agent, but acts indirectly. We noticed that 
dry calico was but little affected by chlorine. The reason 
for this is that chlorine in its great chemism for hydrogen 
abstracts it from the water, and the nascent oxygen unites 
with the coloring matter of the cloth, converting it into 
colorless compounds ; whereas in the dry cloth there was 
comparatively little moisture to furnish the necessary 
oxygen. 

23. Its most extensive use in manufactures is in bleach- 
ing cotton and linen goods and paper pulp. Here, how- 
ever, it is used in the form of bleaching powder. This is 
a compound, which when treated with dilute acid readily 
gives up its chlorine. The following diagram will illus- 
trate the method employed in bleaching cloth. 




Fig. 33. — Cloth-bleaching Apparatus. 

The cloth is seen in a roll at A ; from here it passes 
down under rollers at the bottom of the vat B, which con- 
tains bleaching powder in water, next up over rollers and 
down into a second vat containing dilute hydrochloric acid, 
into a third vat with bleaching powder, and so on until 
the cloth is sufficiently bleached. The excess of chlorine 



112 MODERN CHEMISTRY 

must now be removed, or it will attack the fibers of the 
cloth and make them weak. To prevent this the cloth is 
drawn through another vat, i>, containing an antichlor ; 
that is, a solution which combines with the chlorine still 
present and forms such compounds as will not attack the 
fibres. For this, sodium hyposulphite is frequently used. 
Then, after passing through a vat of water for washing, 
the cloth comes out pure and white» 

Chlorine is also used to some extent as a disinfectant, 
but generally in the form of bleaching powder for this 
purpose also. 

Hydrochloric Acid, HC1 

24. History. — This acid, sold usually under the name 
muriatic acid, has been known for four centuries, and was 
formerly called spirit of salt. Later it received the name 
of marine acid. 

25. Where found. — It is found uncombined only in 
very small quantities. It is said to exist in the stomach 
and to aid digestion, and is sometimes emitted from vol- 
canoes in eruption. 

26. How to prepare Hydrochloric Acid. — The method 
of preparation has already been suggested in one of the 
experiments for making chlorine. 

Experiment 72. — Into a generating flask put about 2 g. of sodium 
chloride, NaCI, and cover with moderately dilute sulphuric acid. The 
action will be seen to begin immediately, but it is better to warm 
gently. Collect two or three bottles of the gas by downward displace- 
ment and preserve for a study of the properties. Keep them covered 
to prevent diffusion. The bottle is full when a moistened piece of 
blue litmus paper held near the mouth is quickly turned red. 

27. Manufacture on a Large Scale. — The method illus- 
trated by this experiment is really the one used in prepar- 
ing hydrochloric acid on a large scale. It is nearly all 



THE HALOGENS 



113 



obtained as a by-product in the manufacture of soda crys- 
tals preparatory to the making of soap. Like many other 
valuable articles of commerce, it was at one time allowed 
to go to waste as of no value. 

28. In the manufacture of sodium carbonate, common 
salt was treated with sulphuric acid as above, and the gas 
obtained was allowed to escape from the flues. But being 
heavier than the air it settled to the ground, destroying 
vegetation and rendering all life in the neighborhood 
almost unendurable. In some places it was produced so 
abundantly as to corrode even the tools of workmen. It 
thus became so great an evil that laws were passed pro- 
hibiting any manufacturer from allowing the escape of 
such gas, just as the consumption of coal smoke is de- 
manded in most large cities to-day. An attempt was also 
made to conduct the gases into streams of water, but this 
resulted in the death of animals living in the streams. 

29. Finally uses were found for the acid, and then plans 
were thought of and efforts made to save and use it. The 
gas is conducted into 
towers filled loosely 
with coke, down which 
water is allowed to 
trickle slowly. In this 
way the gas is practi- 
cally all absorbed, and 
there results a moder- 
ately strong aqueous 
solution of hydrochlo- 
ric acid. Sometimes 
the gases are conducted 
through large Woulff bottles partly filled with water, where 
solution is effected in the same way. 




Fig. 34. —Hydrochloric Acid Factory. 



114 



MODERN CHEMISTRY 



The reaction that takes place may be represented as 
follows : — 

NaCl + H 2 S0 4 = NaHS0 4 + HC1. 

If, however, the heat is increased, a larger amount of 
hydrochloric acid is obtained by using the same amount 
of sulphuric acid with more salt. Thus : — 

2 NaCl + H 2 S0 4 = Na 2 S0 4 + 2 HC1. 

30. Experiments with Hydrochloric Acid. — Many char- 
acteristics of hydrochloric acid may be learned by the 
following experiments : — 

Experiment 73. — Into a bottle of the gas collected above put 
moistened pieces of blue and red litmus paper. How are they affected ? 
Lower a candle into the bottle. What happens? Will the gas burn ? 
Experiment 74. — To show the solubility of the gas in water. 
Add to a bottle of hydrochloric acid gas a little water and shake for a 
moment. Hold a piece of moistened blue litmus paper within the bottle. 
Is it affected? Drop it into the solution. What happens? What has 
the water done ? Has the solution any taste ? 
Experiment 75. — Purpose same as the 
preceding. This is a repetition of the 
"ammonia f ountain " experiment. In pre- 
paring for it one or two additional points 
should be noticed. It is better to use appa- 
ratus somewhat smaller than before, and the 
gas must be collected by downward displace- 
ment. The lower flask in this case had bet- 
ter be fitted with a two-hole rubber cork, 
through one of which the long tube passes. 
Through the other should be passed a short 
tube bent at right angles, as shown in the 
figure accompanying. 

When the flask is well filled with gas, 
make the connections all tight, then blow 
Fig. 35. —Hydrochloric trough the bent tube b to start the flow. 
Acid Fountain. Otherwise it will be necessary to wait several 




THE HALOGENS 115 

minutes before the water will enter the upper flask. The experiment 
works well, but will be more attractive if the water is colored by lit- 
mus or some vegetable solution which will change color upon absorbing 
the acid in the upper flask. A drop or two of ammonia and a few of 
phenol phthalein in the water serve excellently. The deep purplish 
red solution becomes perfectly colorless as it enters the upper flask. 

31 . Characteristics of Hydrochloric Acid. — Hydrochloric 
acid is a colorless gas, somewhat heavier than air, and has 
a very irritating odor. It neither burns nor supports com- 
bustion ; it turns blue litmus paper red, and is very soluble 
in water. At 0° C. 1 liter of water will dissolve about 500 
liters of hydrochloric acid gas. So great is its affinity for 
moisture that whenever it escapes into damp air, heavy, 
white clouds appear. 

32. The commercial acid, which is simply an aqueous 
solution of the gas, contains about 32 per cent of acid. 
Very dilute solutions of hydrochloric acid may be concen- 
trated by heating until the solution contains 20 per cent 
of acid, but the process can be carried no further. On 
the other hand, very strong acid, if exposed to the air, or 
if heated, loses strength. 

33. Hydrochloric acid has great affinity for ammonia ; 
if a bottle of hydrochloric acid and a bottle of ammonia 
remain undisturbed side hj side for some time, they 
become thickly coated about the top with ammonium 
chloride, a white salt formed by the union of the two 
gases. 

34. Uses of Hydrochloric Acid. — The chief use of this 
acid is in the preparation of chlorine for the manufacture 
of bleaching powder. It is also used very largely in all 
chemical laboratories as a reagent, in gas works to neu- 
tralize the ammonia solutions drawn off from the "washer," 
and in the preparation of various chlorides. 



116 MODERN CHEMISTRY 

Bromine : Br == 80 

35. Where found. — Because of its great chemical activ- 
ity, bromine, like chlorine, does not occur free, but is 
found in sea water and in salt wells combined with other 
substances. Its discovery dates from the year 1826, when 
Balard found it in sea water. 

36. Commercial Supply. — The greater part of the com- 
mercial supply of bromine is obtained from Germany and 
the United States. The greater amount used in this 
country comes from Pomeroy Bend, Ohio, where there are 
a large number of salt wells. Bromine appears there in 
the form of magnesium and sodium bromide. The salt 
water from these wells is boiled down to a certain extent, 
the common salt (NaCl) crystallizing out, while the other 
compounds remain in solution. This residue is known as 
the "mother liquor." The next step in the process is to 
/put the solution into stills hewn out of solid rock, adding 
to it manganese dioxide and sulphuric acid. The whole 
is then heated by steam introduced into the liquid through 
pipes. Bromine distills over and is condensed under water. 

37. Formerly bromine was expensive, but, owing to 
cheaper methods of production, the price has been so re- 
duced that many of the salt works no longer prepare it. 
The method of preparation described above is illustrated 
in the following experiment : — 

Experiment 76. — Into a test-tube put a few small crystals of 
sodium or potassium bromide, add a little manganese dioxide, and 
cover with sulphuric acid. Warm slightly and notice the dark red 
gas given off. What other gas have we prepared that resembles this 
somewhat? Describe the odor. How does it affect the eyes? Try 
its bleaching effects upon a moistened piece of calico or litmus paper. 
How does it compare with chlorine in this respect ? Does anything 
condense upon the cooler portion of the tube ? What is its physical 
condition? Its color? 



TEE HALOGENS 117 

38. Laboratory Method of obtaining Bromine. — If some 
bromine is desired for class experiments, it may be pre- 
pared as above. Attach a delivery tube and conduct the 
gas into cold water. As soon as the water is saturated, 
the bromine will condense in the bottom of the jar. It 
may then be obtained from the water by a separating 
funnel, or by pouring into a burette and drawing off the 
heavier liquid as needed for experiment. Preserve both 
the bromine and the water. 

39. The reaction that takes place is the same as in the 
preparation of chlorine by a similar method, thus : — 

2 KBr + 2 H 2 S0 4 + Mn0 2 = 

K 2 S0 4 + MnS0 4 + 2 H 2 + Br 2 . 

40. Another Method. — Sometimes another method is 
used when the purpose is merely to determine whether 
bromine is present in a solution. In this process the 
bromine is set free from its compound by the use of 
chlorine. 

Experiment 77. — To the solution supposed to contain bromine 
add a little chlorine water as prepared in Experiment 61. If bromine 
is present, the solution should turn darker in color, due to the libera- 
tion of the bromine by the chlorine. Which does this experiment 
show to have the greater chemism ? To prove that this color is due 
to the presence of free bromine add about a half cubic centimeter of 
carbon disulphide, shake well, and allow it to settle. If free bromine 
is present, the disulphide will be turned brown from the fact that it 
has taken up all the free bromine in the solution. 

41. Characteristics of Bromine. — Bromine is a dark 
reddish brown liquid. It is the only non-metallic ele- 
ment that is a liquid. It is very volatile, giving off at 
all temperatures heavy brown fumes. At seven degrees 
below zero it solidifies. It has a very disagreeable odor, 



118 MODERN CHEMISTRY 

and attacks not only the throat and nostrils, but also the 
eyes. It differs from chlorine in that the odor is more 
sickening, and it was this fact that gave to the element 
the Greek name brornos, meaning offensive odor. 

42. The vapors are non-combustible, yet, like chlorine, 
they allow of the continued combustion of a jet of hydro- 
gen. As the hydrogen burns, the red vapors gradually 
disappear, and colorless hydrobromic acid gas takes their 
place. Powdered arsenic, sifted into the vapors, burns, 
and a small bit of antimony dropped upon liquid bromine 
burns brightly, and the heat generated by the chemical 
action melts the metal, which spins around upon the sur- 
face like sodium upon water. Bromine is soluble to a 
considerable extent in water, and if the temperature of 
such a solution is reduced by surrounding it with a freez- 
ing mixture, light brown crystals of bromine hydroxide 
separate, as did the crystals of chlorine hydroxide under 
similar circumstances. 

43. Experiments with Bromine. — Let the teacher prove 
the above facts by experiments with bromine before the 
class. 

Experiment 78. — Place a small piece of phosphorus in a defla- 
grating spoon and put it into a jar of bromine vapor. Allow it to 
remain a few minutes. Does it burn? Compare bromine with chlo- 
rine in this regard. 

Experiment 79. — To test the bleaching effects of bromine upon 
colored solutions. Pour into a bottle a little bromine vapor, and add 
a few cubic centimeters of logwood, litmus, or carmine solution. Shake 
it up. Notice the effect upon the color. 

44. Uses of Bromine. — The principal use of bromine is 
as a disinfectant. It is also used in organic work in chem- 
istry and in the preparation of some dyes. For organic 
colors it is a strong bleaching agent, though not as active 



THE HALOGENS 119 

as chlorine. There are also several compounds which 
have application in medicine. Of these magnesium bro- 
mide, MgBr 2 , and potassium bromide, KBr, are the most 
important; the former is found in the water of many 
mineral springs and is regarded as of medicinal value ; 
the latter is used as a sedative in the case of nervous 
headache. A third compound, silver bromide, AgBr, is 
used in photography for sensitizing various printing 

papers. 

Iodine : I = 127 

45. The Source of Supply. — Until within recent years, 
the iodine of commerce was obtained from certain varieties 
of sea-weeds. These weeds were collected in large quan- 
tities and burned, and the ashes treated with water to 
dissolve out the sodium carbonate which was wanted for 
making soap. If such sea- weeds are burned at a low tem- 
perature, the iodine will remain in the ashes in the form 
of sodium and potassium iodide. From these it can be 
obtained as shown below. 

46. The greater part of our present supply of iodine 
comes from Chile. There is a desert in that country 
many square miles in area, where are found vast deposits 
of sodium nitrate mixed with considerable quantities of 
soil and small amounts of iodine compounds. This mix- 
ture is treated with water, which dissolves out the sodium 
nitrate and the iodides ; the solution is then evaporated 
till the sodium nitrate crystallizes out, as in the manufac- 
ture of bromine in Ohio, leaving the iodine compounds 
still in solution. The residual solution, called the " mother 
liquor," is treated with manganese dioxide and sulphuric 
acid, and gently heated. 

47. Preparation for Commerce. — When treated as above, 
from the mother liquor, violet fumes of vaporous iodine 



120 



MODERN CHEMISTRY 



are given off abundantly ; they are passed over into cool 
chambers, where they condense. To further purify the 
iodine, it is resublimed at a low temperature and con- 
densed in a series of conical-shaped flasks (see Fig. 36). 




Fig. 36. — Iodine Apparatus. 



At the left is a small brick furnace, in the upper part of 
which is an oven. The iodine to be purified is placed in 
the oven, and gently heated. The final reaction in the 
separation of the iodine is the same as in the case of the 
bromine and chlorine. 

2 NaI + 2 H 2 S0 4 + Mn0 2 = Na a S0 4 + MnS0 4 + 2 H 2 + I 2 . 

The essential features of this method of preparing iodine 
may be shown by the following experiment : — 

Experiment 80. — Into a small test-tube put a crystal or two of 
potassium iodide, add a little manganese dioxide, and cover with sul- 
phuric acid. Warm gently ; notice the fumes that are given off: and 
what condenses upon the cooler portion of the tube. 

48. Another Method of preparing Iodine. — The follow- 
ing method is employed to some extent in France in 
obtaining iodine from the ashes of sea-weeds. It is also 
the usual method pursued in the laboratory in testing for 
iodine. The plan consists simply in adding free chlorine 






THE HALOGENS 121 

to the iodine solution, whereby the latter is liberated from 
its compounds. As a commercial method it is open to the 
objection that if too little chlorine is added, not all of 
the iodine is liberated, and if too much, a portion is 
redissolved. 

Experiment 81. — To a solution containing iodine in combina- 
tion add a few drops of chlorine water. What change in color takes 
place ? This indicates free iodine, as maybe proven by adding starch 
paste solution. The starch will turn blue, as it did with ozone. 

49. Experiments with Iodine. — Many characteristics of 
iodine may be learned from the following experiment : — 

Experiment 82. — Put a small crystal of iodine into a test-tube 
and warm gently. What happens? Describe the color and odor of 
the vapors. Hold a piece of moistened starch paper near the mouth 
of the test-tube ; how is the starch affected ? Close the mouth of the 
tube with your finger and notice the stain that is formed. See whether 
you can remove it by moistening with caustic potash or ammonia. 

50. Characteristics of Iodine. — Iodine is a solid of a 
dark bluish black color, with a metallic luster. At or- 
dinary temperatures it is somewhat volatile, and when 
gently heated it is readily converted into vapors of a 
beautiful violet color. It was from this fact that the 
element received its name, iodine being derived from a 
Greek word which means violet. The odor of the vapors 
resembles somewhat that of dilute chlorine, but is less 
irritating. It has the power of turning the skin yellow, 
but the stain may be removed by treatment with some 
alkali. It has feeble bleaching properties, and turns 
starch paste blue. This is so delicate a test that one 
part of iodine in several hundred thousand of water will 
be clearly shown. Its affinity for phosphorus is so strong 
that if a crystal of iodine be dropped upon a small piece of 
phosphorus, the latter will be ignited almost instantly. 



122 MODERN CHEMISTRY 

51. Solvents for Iodine. — Among the better solvents 
for iodine are chloroform, ether, alcohol, carbon disul- 
phide, and a solution of potassium iodide. 

Experiment 83. — Put a crystal of iodine into a test-tube with a 
little cold water. Shake for a moment or two, and then pour off a 
part of it into another tube and test with starch paste to determine 
whether any has dissolved. What are your conclusions ? Warm the 
remainder; what indications are there that the iodine is dissolving? 
Test the solution again with starch paste, or with carbon disulphide, 
thus : add about a half cubic centimeter of the disulphide to the iodine 
solution, and shake well. Notice the beautiful violet color imparted 
to the disulphide. 

Try alcohol also as a solvent. Before testing the solution with 
starch or carbon disulphide, dilute until pale yellow in color. What 
are the results ? Try in the same way a solution of potassium iodide 
upon an iodine crystal, and state results. 

52. Uses for Iodine. — In the form of a tincture, or 
alcoholic solution, iodine is used to a considerable extent 
in medicine to prevent the spread of eruptive diseases, like 
erysipelas, in skin affections, sore throat, and the like. In 
the compound iodoform it is used by physicians as a deo- 
dorant and disinfectant. As potassium or sodium iodide 
it is frequently used as a reagent in the laboratory, and 
to a limited extent in making aniline dyes. In these vari- 
ous ways 300 tons or more are used annually. 

53. Some Comparisons. — It has probably been observed 
that the same method is used in preparing chlorine, bro- 
mine, and iodine. Notice the following reactions : — 

CI Mn0 2 + 2 H 2 S0 4 + 2NaCl = MnS0 4 + Na 2 S0 4 
+ 2 H 2 + Cl v 
• Mn0 2 + 2H 2 S0 4 +2i^£r = MnS0 4 +Na 2 S0 4 
+ 2 H 2 + Br 2 . 
Mn0 9 + 2 H 9 S0 4 + 2 MI * MnS0 4 + Na 2 S0 4 
J + 2H 2 0+7 2 , 



THE HALOGENS 123 

SUMMARY OF CHAPTER 

Meaning of term halogen. 

Names, symbols, and atomic weights of the halogens. 

Comparison of the halogens. 

a. Method of preparing. 

b. Physical condition at ordinary temperatures ; at lower tem- 

peratures. 

c. Color. 

d. Odor. 

e. Density. 

/. Chemical activity. 

g. Bleaching powers. 

h. Affinity for certain substances, as hydrogen, phosphorus, etc. 

i. Uses. 

/. Hydrogen compounds. 

Compare hydrofluoric and hydrochloric acids as to — 

1. Method of preparation. 

2. Characteristics. 

3. Uses. 
Special points for study. 

Method of etching glass. 

What kind of substances may be used instead of manganese 
dioxide in preparing chlorine ? Why ? 
Proof of this by experiments. 
Practical application of this. 
Compare these two methods of making chlorine : — 

1. Mn0 2 + HCl. 

2. MnO, + XaCl + H 2 S0 4 . 

How similar? How different? 
Explain haw chlorine bleaches. Write the equation. Source of 

commercial supply of hydrochloric acid. 
Tests for bromine and iodine with carbon disulphide — compare 

results. 
Method of obtaining and purifying iodine. 
Describe experiments which illustrate chief properties of chlorine, 

bromine, and iodine. 
Solvents for chlorine, bromine, and iodine. 



CHAPTER X 
ACIDS, ALKALIES, AND SALTS 

1. Neutralization. — There are a great many substances 
which, if put together, have the power of destroying the 
characteristic properties of each other. 

Experiment 84. — To show this fact, put into an evaporating dish 
about 10 cc. of dilute hydrochloric acid and dip into it a small piece 
of blue litmus paper. Notice that it is changed to red. Now add 
slowly, stirring all the time with a glass rod, a solution of caustic 
soda, until the litmus paper just turns blue again ; then add one drop 
of hydrochloric acid. You ought now to have a solution that will not 
affect either red or blue litmus paper. Boil this solution to dryness. 
What is the appearance of the solid thus obtained? Taste it. Does 
it seem familiar ? Dissolve it in a little water and test with both red 
and blue litmus paper. Is the paper affected? Now boil a little 
hydrochloric acid to dryness. Does it leave a residue? Examine a 
specimen of solid caustic soda and compare with the white solid 
obtained above. Are the two solids the same? Do they both affect 
litmus in the same way? 

2. From this experiment we see that the acid and the 
caustic soda, on being put together, have both lost their 
characteristic properties and have reacted to form a new 
substance having the properties of neither. In other 
words, they have neutralized each other. 

3. Bases. — Such substances as have the power of neu- 
tralizing the properties of acids are called bases. This was 
shown in Experiment 84 above. We have already seen 
that the compound of any element with oxygen is called 
an oxide; many of the oxides combine with water to form 

124 



ACIDS, ALKALIES, AND SALTS 125 

water oxides, or, to use the ordinary term, which is from 
the Greek, hydroxides or hydrates. We have seen also 
that some oxides or anhydrides, when united with water, 
form acids, as, for example, nitrogen trioxide. Strictly 
speaking, such compounds are hydroxides, but we never 
apply that term to them; it is restricted altogether to 
those compounds of oxides with water which have the 
power of neutralizing acids. 

4. Alkalies. — A few of these compounds, which have 
exceedingly strong basic properties, are called alkalies. 
The four most common alkalies are the three lrydroxides 
of sodium, potassium and calcium, and ammonia. If we 
study the formulae of the hydroxides, we shall see that 
water may be taken as the type upon which all the others 
are built. Thus : — 

Water ........ HOH 

Caustic Potash KOH 

Caustic Soda NaOH 

Lime Water Ca(OH) 2 

Ammonium Hydroxide . . NH 4 OH 

The only difference is that one atom of hydrogen in the 
water has been replaced by some metal or group of elements. 

5. As most of the metals themselves have all the char- 
acteristic properties of bases, they are, by many, regarded 
as bases, and always spoken of as such. Probably there 
is no serious objection to this, but it should be remembered 
that, strictly speaking, all bases are compound bodies. 

6. Acids. — It is a difficult matter to define acids. They 
are substances which have properties exactly the opposite 
of bases. They possess the power, not only of turning 



126 MODERN CHEMISTRY 

blue litmus red, but of similarly affecting various other 
vegetable colors, all of which are restored again by the 
use of an alkali. They also have a sour taste, though 
this is not a distinctive feature, as many bodies not acids 
also have the same property. 

7. Their Composition. — If we recall the formulae of the 
few acids we have already used, nitric, hydrochloric, and 
sulphuric, we see that they all contain hydrogen ; it is 
true also that most contain oxygen, together with some 
third element which seems to give the distinctive proper- 
ties to the acid. It was at one time supposed that all 
acids contained oxygen, and in accordance with this idea 
oxygen received its name. Later, however, were discov- 
ered hydrochloric and other acids, which contained no 
oxygen whatever. A distinctive property of acids is that 
they all have the power of giving up the whole or a part 
of their hydrogen, and of combining instead with some 
metal or base. Thus we have seen that when zinc was 
treated with hydrochloric or sulphuric acid, the metal 
replaced the hydrogen in the acid, setting the former free. 
(See p. 40.) 

8. Salts. — A salt is a compound formed by the union 
of an acid with a base or metal, possessing properties 
different from those of either of its constituents. We 
have been accustomed to think of salt as a particular sub- 
stance used in seasoning food, but we must now remember 
that it is a term applied to a large number of compounds, 
called salts because they resemble common salt in appear- 
ance or properties. They are all produced in the same 
way. We saw above that when we neutralized hydro- 
chloric acid with caustic soda and boiled to dryness, we 
obtained a white solid, resembling and tasting like com- 
mon salt, which it really was. 






ACIDS, ALKALIES, AND SALTS 127 

Experiment 85. — In the same way as in Experiment 84, neu- 
tralize about 10 cc. of hydrochloric acid with caustic potash and boil 
to dryness as before. Compare the salt produced, in taste and appear- 
ance, with that obtained before. 

9. Normal or Neutral Salts. — There are two general 
classes of salts, neutral or normal, and acid. A normal 
salt is one in which all the displaceable hydrogen of the 
acid used in making the salt has been replaced by some 
base. For example, when caustic potash and sulphuric 
acid neutralize one another, the following reaction takes 
place : — 

H 2 S0 4 + 2 KOH = K 2 S0 4 + 2 H 2 0. 

We see here that all the hydrogen in the sulphuric acid, 
two atoms, has been replaced by an equivalent amount of 
the metal, potassium, and the salt produced, potassium 
sulphate, K 2 S0 4 , is a normal salt. 

10. Again, if lead is treated with vinegar (acetic acid), 
which is represented by the formula HC 2 H 3 2 , we have 
this reaction : — 

Fb + 2 HC 2 H 3 2 = Pb(C 2 H 3 2 ) 2 + H 2 . 

It will be noticed that in the salt resulting, Pb(C 2 H 3 2 ) 2 , 
a quantity of hydrogen remains. Lead acetate is, notwith- 
standing, a neutral salt, because only one atom of hydro- 
gen, the first, in each molecule of acid can be displaced. 

11. Acid Salts. — If, however, we use only half the 
amount of caustic potash shown by the first reaction above 
with the sulphuric acid, we shall replace only half of the 
hydrogen in the acid, and the salt resulting will be an 
acid salt, thus : — 

H 2 S0 4 + KOH = KHS0 4 + H 2 0. 



128 MODERN CHEMISTRY 

12. Reading the Formulae of Salts. — The compound, 
K 2 S0 4 , is read, normal potassium sulphate, or usually, 
simply potassium sulphate. The acid salt, KHS0 4 , is read, 
acid potassium sulphate, or potassium hydrogen sulphate. 
Sometimes the prefix mono is applied, but usually only in 
the case of salts of acids having three or more replaceable 
hydrogen atoms, as phosphoric, H 3 P0 4 , or silicic, H 4 Si0 4 . 
With these acids we may form the following salts : — 

Phosphoric Acid, H 3 P0 4 

Mono-sodium Phosphate NaH 2 P0 4 

Di-sodium Phosphate Na 2 HP0 4 

Normal sodium Phosphate Na 3 P0 4 

Silicic Acid, H 4 Si0 4 

Mono-sodium Silicate . NaH 3 Si0 4 

Di-sodium Silicate Na 2 H 2 Si0 4 

Tri-sodium Silicate Na 3 HSi0 4 

Normal sodium Silicate Na 4 Si0 4 

Exercises. — In the following formulae, state which represent 
acids, which bases, and which salts, giving reasons therefor. Give 
also the name of the substance represented. If salt, state whether 
acid or normal : — 

JSTa 2 S0 4 , KOH, H 2 S0 4 , H 3 P0 3 , ZnS0 4 , KN0 3 , Ca(0H) 9 , BaS0 4 , 
K 3 P0 4 , K 2 HP0 4 , KH 2 P0 4 , HC1, NaOH, NaHS0 4 , NaN0 3 ." 

13. Nomenclature of Acids. — It will be noticed that 
with one exception the acids we have met with so far all 
have names ending in ic ; thus : — 

Sulphuric .... H 2 S0 4 

Nitric HN0 3 

Phosphoric . . . . H 3 P0 4 

Silicic H 4 Si0 4 , etc. 



ACIDS, ALKALIES, AND SALTS 



129 



The greater number of acids with which we shall have 
to deal, as already stated, contain three elements, the first 
of which is hydrogen, the third oxygen, and a second 
which gives the name to the acid. Thus the middle sym- 
bols in the above formulae are : — 



s . 


Sulphur 


acid, H 2 S0 4 . 


Sulphuric 


N . 


Nitrogen 


HN0 3 . 


. Nitric 


P . 


Phosphorus 


' " H 3 P0 4 . 


Phosphoric 


Si . 


. Silicon 


: "' H 4 SiO,. 


Silicic 



14. Sometimes, however, this second element forms 
more than one acid with hydrogen and oxygen. In such 
cases the most common, and hence the earliest knoAvn, 
received the name with the termination ie. Then the 
acid having a smaller amount of oxygen is given the 
termination ous. This we have seen in the two nitro- 
gen acids : — 



Nitric . 
Nitrous 



HN0 3 
HN0 



Oxygen, 3 atoms. 



15. Sometimes even three or four acids are formed from 
the same three elements, the amount of oxygen only vary- 
ing. In such cases, the one with the least quantity of 
oxygen is given the prefix hypo, meaning under or lesser, 
and the one with the most oxygen has the prefix per, 
beyond or above. These may be illustrated by the follow- 
ing series : — 



Sulphune . 


• H 2 


S 


o 4 • 


Oxygen, 4 atoms 


Sulphurous . . . 


. H 2 


S 


3 . 


3 " 


#y/>0-sulphurous . 


. H 2 


s 


o 2 . 


" 2 " 



130 MODERN CHEMISTRY 

Chlorine Acids 



Perchloric 


. H 


CI 


o 4 • 


. Oxygen, 4 atoms. 


Chloric . 


. H 


CI 


O3 • 


3 « 


Chlorous 


. H 


CI 


2 . 


" 2 " 


Hypo-chlorous . 


. H 


CI 


. 


" 1 " 



16. Nomenclature of Salts. — All of the acids named 
above have the power of combining with various metals, 
or their hydroxides, to form salts. All such as result 
from the union of a base with an ic acid are given names 
with the termination ate. Thus, all salts of sulphuric 
acid are sulphates; of nitric acid, nitrates; phosphoric acid, 
phosphates, etc. To illustrate : — 

H 2 S0 4 , sulphuric acid, gives — 

with zinc, ZnS0 4 , zinc sulphate ; 

with potassium, K 2 S0 4 , potassium sulphate ; 

with calcium, CaS0 4 , calcium sulphate. 

HN0 3 , nitric acid, gives — 

with potassium, KN0 3 , potassium nitrate ; 

with sodium, NaN0 3 , sodium nitrate; 

with ammonia, NH 4 N0 3 , ammonium nitrate. 

17. Salts formed from the ons acids receive names end- 
ing in ite. (It may aid the memory in associating the 
pronouns singular, _T, plural objective us, with the ous 
acids and ite salts.) Thus, from 

H 2 S0 3 , sulphurous acid, we have 
K 2 S0 3 , potassium sulphite ; 
Na 2 S0 3 , sodium sulphite, etc. 



ACIDS, ALKALIES, AND SALTS 



131 



In the case of salts formed from the hypo and per 
acids, the corresponding prefix is simply given to the 
salt. Thus : — 

NaCIO, sodium hypochlorite, from 
HCIO, acid, hypochlorous, and 
NaC10 4 , sodium perchlorate, from 
HC10 4 , acid, perchloric. 

18. Binary Compounds. — All of the above salts are 
formed from what are sometimes called ternary acids; 
that is, those consisting of three (or more) terms. In 
like manner, a binary compound would be one which 
consists of only tw r o elements. The following are ex- 
amples : — 

Common Salt .... NaCl 
Calcium Chloride . . . CaCl 2 
Water H 2 



Turpentine ^ 10 H 



16 



It will be noticed from these formulae that though in a 
binary compound there are but two elements, the number 
of atoms of each of these elements is quite variable. 

19. As already stated, there are a few acids which 
contain no oxygen. Salts obtained from them would, 
therefore, all be binaries. Thus : — 

from Hydrochloric acid, HC1, we obtain the chlorides ; 
Hydrobromic " HBr, " bromides; 

Hydriodic " HI, " iodides; 

Hydrofluoric " HF, " fluorides; 

Hydrosulphuric " H 2 S, " sulphides. 



132 



MODERN CHEMISTRY 



20. Compounds with Oxygen Alone. — Oxides. — It will 
be noticed that all binary salts are given names ending in 
ide. Furthermore, it is seen that it is the second element 
in every case which gives the name to the substance-; 
thus : — 

NaCl 



KC1 
MnCl 2 

CaCL 



► are all chlorides, while the first 



element indicated in the formula is simply descriptive in 
character, or the adjective that tells what kind of a 
chloride. Thus the above are 



Sodium 
Potassium 
Manganese 
Calcium 



Chloride ; 



just as we might say 



Stone 
Brick 
Frame 
Log 



House. 



21. It frequently happens that two elements, just as in 
the case of the ternary acids, may unite in different pro- 
portions to form two or more compounds. We have 
already seen this in studying the oxides of nitrogen, p. 81. 
When two such exist, as for example the two oxides 
of mercury, HgO and Hg 2 0, the one having the smaller 
proportion of the second element, as indicated by the 
formula, is the ous compound, just as in the case of the 
acids already studied, while the one having the greater 
proportion of the same element is the ic compound. 



ACIDS, ALKALIES, AND SALTS 133 

22. Again, we noticed in studying the oxides of ni- 
trogen : — 

N 2 0, ratio of oxygen to nitrogen, 1 : 2 

N 2 2 , " " " 2:2 = 1:1. 

In the first we found one-half as many atoms of oxygen as 
of nitrogen ; in the second the same number ; they were 
therefore called nitrous and nitric oxides. In a few 
instances, instead of using the English name with 
the terminations ous and ic, for the sake of euphony, 
the Latin forms are taken. Thus : — 

Cu 2 0, Cuprous Oxide 
CuO, Cupric " 
FeCl 2 , Ferrous Chloride 
Fe 2 Cl 6 , Ferric 

23. Returning to the series of nitrogen compounds, it 
will be noticed that they were given two names. This is 
very often done, one of them using a prefix to indicate 
the exact number of atoms of the last element in the 
formula of the compound. Thus we have 

N 2 0, Nitrogen Monoxide. 
N 2 4 , " Tetroxide 
P 2 5 , Phosphorus Pentoxide, etc. 

24. Old Forms. — Occasionally we use the old terms, 
pro, per, and sesqui. The first is a prefix, meaning before, 
.and is given to some uncommon or unstable compounds 
1 which in the case of a series would be the first or lowest. 

Thus, FeO is sometimes spoken of as iron protoxide. 
■ ] Nitrogen tetroxide, N 2 4 , is also called peroxide, as is 



134 MODERN CHEMISTRY 

hydrogen dioxide as well, it being the compound com- 
ing in the series beyond the others. Sesqui is applied 
to binary compounds in which the two elements unite 
in the ratio of 2 to 3, as in Fe 2 3 , iron sesquioxide. 

SUMMARY OF CHAPTER 

Neutralization — Meaning of the term. 

Experiments to illustrate. 
Three classes of compounds. 
Compare bases and acids. 

a. In composition. 

b. In properties. 
Alkalies — What are they ? 

Examples. 
Salts — What are they ? 
Two classes. 

How formed. 

How distinguished by name. 

Examples to illustrate. 
Nomenclature. 

a. Of acids. 

b. Of salts. 
Examples to illustrate both. 

Binary compounds. 

Meaning of term — Illustrations. 
Six important classes — Examples. 
Nomenclature — Compare with acids. 



CHAPTER XI 

CARBON AND A FEW COMPOUNDS. C = 12 

1. Abundance. — With the exception of oxygen, carbon 
is the most widely distributed and most abundant of all 
the elements. In the form of compounds it is found in 
the air as carbon dioxide, resulting from combustion and 
respiration, and in limestone, CaC0 3 , which constitutes a 
large portion of the rocky crust of the earth. It also 
occurs in almost all food products, such as sugar, flour, 
starch, vegetables, and fruits, and forms a large part of 
the woody structure of plants and trees. 

2. Forms. — In the free state carbon may be considered 
under two divisions : — 

a. Crystallized, including 

1. The Diamond. 

2. Graphite or Plumbago. 

b. Amorphous (without crystalline form), 

1. Coal. 

2. Lampblack. 

3. Gas Carbon, etc. 

3. Diamonds. — The diamond occurs in octahedral crys- 
tals. It is found in South America, Africa, Australia, and 
India. By some the stones are thought to be of meteoric 
origin and not native to the earth, but the theory seems 
not well founded. Moissan, the French chemist, has suc- 
ceeded in making a few diamonds in the electrical furnace, 

135 



136 MODERN CHEMISTRY 

but they have all been exceedingly small, and black in 
color, so as to have no value except in a scientific way. 
In nature they occur rough and covered with a layer of 
partially decomposed rock. The most highly prized are 
perfectly transparent, but many of various colors have 
been found. The diamond has strong refractive power, 
is the hardest of known substances, and can be cut and 
polished only by its own dust. 

4. Their Practical Uses. — Diamonds are used, not only 
as ornaments, but also in cutting glass ; and the cheaper, 
imperfect varieties are employed as tips on drills for cut- 
ting through hard rocks. That the diamond consists of 
carbon may be proved by burning it between electric ter- 
minals in an atmosphere of oxygen ; the diamond and 
oxygen disappear, and carbon dioxide, C0 2 , remains. 

5. Graphite. — Next to the diamond, graphite or plum- 
bago is the purest form of carbon. It is sometimes called 
black lead, but it contains no lead whatever. It is often 
found in hexagonal prisms, is steel-gray in color, has a 
greasy feeling, and as a mass is comparatively soft, though 
the particles themselves are very hard. 

6. That it consists of carbon may be proved by testing 
it in the electric furnace, as in the case of the diamond, 
similar results being obtained. 

7. Uses. — The most common use of graphite is ii| 
making what are known as lead pencils, so named because 
plumbago was at first supposed to be a compound of lead. 
In making pencils the graphite is thoroughly pulverized 
and mixed, according to the grade of pencil, with different 
proportions of fine clay, also well ground. The whole is 
then made up with water into a dough and pressed into 
moulds and dried, or while still soft is forced through 
plates with apertures the size of the lead in the pencil. 



CARBON AND A FEW COMPOUNDS 137 

Graphite is also used as a lubricant, as a stove polish, and 
in making crucibles. 

8. Amorphous Carbon. — The most important uncrys- 
tallized forms of carbon are the various coals, — anthra- 
cite, semi-anthracite, bituminous, lignite, peat, jet, cannel, 
and the artificial form, charcoal. Of great importance 
also are gas carbon, lampblack, and coke. 

9. Coal. — Coal is supposed to be the result of pressure 
and heat applied to a luxuriant vegetable growth in the 
presence of moisture. Peat is the newest of the coals, 
being in process of formation in swamp-lands to-day. It 
consists almost entirely of a mass of roots. Next in age 
is lignite, in which the woody structure is still apparent. 

10. Anthracite and Bituminous Coals. — Anthracite dif- 
fers from bituminous coal in that the former, being sub- 
jected to greater heat and pressure, has been deprived of 
its volatile products. These furnish in part, at least, the 
petroleum and natural gas of the present time. Petroleum 
is really a mixture of a number of different oils, with boil- 
ing points differing greatly. These, in the process of re- 
fining the crude oil, distill over at different temperatures. 
Such light oils as naphtha and benzine are obtained at a 
low temperature, a somewhat higher temperature produc- 
ing kerosene, and higher still paraffine. This method of 
separating substances through differences in their boiling 
points is called fractional distillation, while that in which 
the substance heated is decomposed is called destructive 
distillation. 

11. Charcoal. — Charcoal, because of the abundance of 
timber, has usually been prepared in a simple, but very 
wasteful, manner. Large piles of wood are covered with 
earth and set on fire. Most of the air is excluded in this 
way, and only enough heat is produced to expel the vola- 



^ 



138 



MODERN CHEMISTRY 



tile products from the wood. At present, however, in 
some sections the wood is heated in iron retorts, and the 
volatile products are condensed and refined, much in the 
same way as with petroleum. 

12. Coke. — Coke bears the same relation to soft or 
bituminous coal that charcoal does to wood. It is an 
artificial product obtained by expelling all the volatile 
products from the coal. Most of the supply comes from 
the gas factories as a by-product, but where the local 
supply is insufficient, it is prepared specially for smelters 
in large brick ovens. See the figure below. 




Fig. 37.— Coke Oven. 
aa, openings for slight draught at first. DD, doors for removing coke. 

The coal, in car loads, is shoveled in from above ; it is 
then ignited, and the openings on the side almost entirely 
closed. In the course of several hours the combustion of 
the lower layer of coal has converted the remainder into 
coke, the doors are opened, and the coke drawn out. 

13. Gas Carbon. — Gas carbon is another by-product of 
coal-gas manufacture. Just as soot collects in stove-pipes 
and flues, so on the inside of the retorts there is gradually 
deposited a very hard, black substance, known as gas 
carbon. This is occasionally removed, ground up fine, 



CARBON AND A FEW COMPOUNDS 139 

and moulded into the familiar carbon rods in our electric 
arc lights, and into plates for electric batteries. 

14. Lampblack. — Lampblack is the result of the imper- 
fect combustion of any substance rich in carbon. It is 
usually prepared by burning some hydrocarbon, such as 
turpentine, C 10 H 16 , in a limited supply of air. The dense 
black smoke resulting is allowed to deposit upon canvas 

1 in a cool room, from which it is shaken, and is then ready 
for commerce. It is used in making black paint, printers' 
ink, etc. 

Some Uses of Carbon 

15. As a Reducing Agent. — In the form of charcoal or 
coke, at a high temperature, carbon is a great reducing or 
deoxidizing agent. By this we mean that when it is heated 
with the oxides of various metals, it has the power of 
combining with the oxygen and reducing the oxide to the 
metallic condition. This will be made clear by the fol- 
lowing experiment. 

Experiment 86. — Make a small cavity near one end of a stick of 
charcoal, and put therein a little litharge, PbO, or red lead, Pb 3 4 , 
and heat strongly with the reducing flame. Notice that in a few 

j minutes a bead of lead appears instead of the oxide that we had. 

i The carbon has combined with the oxygen in the lead oxide to form 
carbon dioxide, and the lead has been reduced to the metallic form. 



16. As an Absorbent. — Carbon in the form of charcoal 
is an excellent absorbent, not only of gases, but of certain 
other substances as well. 

Experiment 87. — Thrust a piece of charcoal under water and 
hold it there a minute or so. What is seen escaping from the char- 
I coal? Heat another piece red-hot and plunge under water. Are the 
1 results different? Why?* 

* In this connection refer to Exp. 47, under ammonia. 



140 MODERN CHEMISTRY 

Experiment 88. — Soak some vegetable matter in a vessel of 
water until it has become very offensive, on account of decomposi- 
tion. Put a little of this water into a flask and add some bone-black 
or powdered charcoal, and shake well. Notice that the disagreeable 
odor disappears. 

17. As a Purifier. — • Application is made of this fact in 
purifying cisterns which have become foul with decom- 
posing organic matter. The charcoal should be removed 
after a time and heated to redness to destroy thoroughly 
the organic matter which may have been absorbed. It is 
believed that partial oxidation takes place within the 
pores, but unless the charcoal is heated they eventually 
become clogged. 

Experiment 89. — Fit a filter paper smoothly to a funnel as 
described in Appendix C, page 365, and partly fill it with bone-black. 
Now pour upon it, slowly at first, a few cubic centimeters of logwood 
or some other colored vegetable solution. How is it affected? Try 
also in the same way a solution of copper nitrate. Are the results the 
same ? 

18. In Refining. — An application of the power of char- 
coal to absorb vegetable colors is made in refining sugars. 
At first they are brown, not very different from maple 
sugar in appearance. This raw sugar, as it is called, is 
dissolved in water and passed through filters of bone- 
black which absorb the coloring matter and leave the 
solution clear. This may be shown by filtering a solu- 
tion of molasses in water. 

Experiment 90. — In like manner, charcoal has the power of ab- 
sorbing various organic flavors. Pass through a powdered char- 
coal filter an infusion of tea or coffee, and taste it after it has 
gone through. How is it changed? 

19. It would be impossible to enumerate the various 
uses of carbon in its different forms. Many of these are 



C ABB ON AND A FEW COMPOUNDS 141 

familiar to the student, and others will be learned from 
time to time. Many of them have already been named in 
the sections immediately preceding this. 

Compounds of Carbon. The Oxides 

20. Carbon Monoxide, CO. — This is a gas obtained when 
carbon is burned in a limited supply of air. It may be 
prepared by passing steam over red-hot coke or charcoal, 
whereby the steam is decomposed, thus: — 

H 2 + C = CO + H 2 . 

It is also produced in grates and base-burners. At the 
lower portions of the fire where the heat is most intense 
the carbon is completely burned, producing carbon dioxide ; 
as this passes up through the red-hot coal, it unites with 
another portion of carbon and forms the monoxide. 
Again, on reaching the upper surface, the monoxide unites 
with the oxygen of the air and is burned into carbon 
dioxide. 

21. Carbon monoxide may be prepared in an impure 
form by heating oxalic acid or potassium ferrocyanide with 
sulphuric acid, or by passing a current of carbon dioxide 
slowly through a tube containing red-hot charcoal or coke. 

22. Characteristics of Carbon Monoxide. — Carbon mo- 
noxide is a colorless gas, having a faint, peculiar, but some- 
what unpleasant and stifling odor; it is a little lighter 
than air and burns with a pale blue flame. It is not 
soluble in water, is only slightly explosive when mixed 
with air or oxygen, and is poisonous when inhaled. It 
has the power of decomposing the blood, and thus of ren- 
dering it incapable of carrying oxygen and removing the 
waste of the body. On this account serious results some- 
times follow its escape into rooms from coal stoves when 



142 MODERN CHEMISTRY 

the drafts are closed at night. Open charcoal fires also 
produce the same gas, and have sometimes been the means 
of causing death. 

23. As a Reducing Agent. — It has been seen that carbon 
is a strong reducing agent. Carbon monoxide has the 
same properties, owing to the fact that it has strong 
affinity for more oxygen, to form carbon dioxide. The 
reduction of metallic ores in blast furnaces is, to a con- 
siderable extent, due to this property of carbon monoxide. 
It may be seen by passing a current of carbon monoxide 
over lead oxide, PbO, heated red hot in a tube. The 
monoxide abstracts the oxygen from the lead oxide, form- 
ing carbon dioxide and metallic lead. The reaction is as 
follows: — pb0 + co = pb + C02 

24. Carbon Dioxide, C0 2 . — Where found. — This gas is 
always found in the air, being produced by the combustion 
of organic bodies and by respiration. The proportion 
varies somewhat, but seldom exceeds four parts in 10,000 
parts of air. Another source of this gas as found in the 
atmosphere is fermentation and decay. 

25. Produced in Decomposition. — As already mentioned 
in considering ammonia, organic substances are very un- 
stable and readily break up to form simpler compounds. 
The molecules of most so-called organic compounds consist 
of carbon, hydrogen, oxygen, and often nitrogen, and are 
usually very complicated. In the processes of decay the 
atoms rearrange themselves, and carbon dioxide is one of 
the new products. The process is the same when fer- 
mentation is induced by bacteria or germs, such as those 
of ordinary yeast. If into a flask containing some water j 
sweetened with sugar or molasses a little yeast be intro- , 
duced, fermentation very soon begins, and the bubbles of 



CARBON AND A FEW COMPOUNDS 143 

gas which pass off may be collected and proved to be 
carbon dioxide. 

26. How prepared. — For laboratory purposes carbon 
dioxide is usually prepared by treating some carbonate, as 
marble, CaC0 3 , with dilute acid. 

Experiment 91. — Put into a small flask some marble, coarsely 
powdered, and add some dilute hydrochloric or nitric acid. Xotice 
the rapid effervescence and evolution of colorless gas. 

Experiment 92. — Collect by downward displacement a bottle of 
the gas, generated as above. Lower into it a burning match or candle ; 
what are the results? Ignite a piece of magnesium ribbon and hold 
it in a bottle of carbon dioxide ; what are the results ? What two 
products are formed? Why does the ribbon continue to burn? 
Mention some other gas that supports the combustion of phosphorus, 
but not that of ordinary substances. 

Experiment 93. — To show the density of the gas. Put into a good- 
sized bottle or beaker a small candle and pour in upon it another 
bottle of carbon dioxide. You cannot 
see anything being turned out, but the 
results are apparent. This is sometimes 
, made more effective by fastening at 
short intervals upon the bottom of a 
trough several candles. Lift one end of 
the trough and pour down it a large 
bottle of carbon dioxide. The candles 
will be extinguished, one after another, 
as the gas reaches them. 

Experiment 94. — Purpose same as ~~" p IG 3 g 

preceding. Put into an evaporating dish 

a little gasoline and ignite it. Take a large bottle of carbon dioxide 
and pour suddenly upon the burning oil. The flame will be instantly 
extinguished. 

Experiment 95. — To show effect of carbon dioxide on limestone. 
Pass a current of carbon dioxide through a few cubic centimeters of 
'lime water. Xotice the formation of a white precipitate, which is 
calcium carbonate, CaC0 3 , of the same composition as limestone. 
Continue passing the gas through the milky solution; what change 
takes place? Can you explain ? 




144 MODERN CHEMISTRY 

27. Characteristics of Carbon Dioxide. — From the above 
experiments we learn that carbon dioxide is a colorless, 
odorless gas, considerably heavier than air. It is non- 
combustible and a non-supporter of ordinary combustion, 
though such substances as magnesium, which burns with 
great intensity, are able to decompose the gas and make 
use of the oxygen. It is slightly soluble in water and 
gives to the latter a faint acid taste and reaction. The 
presence of carbon dioxide may always be determined by 
its effect upon lime-water. It forms in the water a white 
precipitate which dissolves slowly again in excess of the 
dioxide. Limestone caves are a manifestation on a large 
scale of the principle shown in the simple experiment 
above. Water under pressure absorbs considerable quan- 
tities of carbon dioxide, which gradually dissolves the 
limestone and forms caverns. 

28. Liquid Carbon Dioxide. — Carbon dioxide may be 
liquefied in strong cylinders by pressure ; if the pressure 
is suddenly withdrawn, a portion of the liquid is rapidly 
vaporized, producing such cold as to convert the remainder 
into a white crystalline solid like snow. The temperature of 
this carbonic acid snow is sufficiently low to freeze mercury. 
The solid carbon dioxide vaporizes without first melting. 

29. Choke Damp. — Because of its density, carbon di- 
oxide frequently collects in deserted mines and deep wells, 
and is called by miners "choke damp." Its presence in 
such places, however, may always be detected by lowering 
to the bottom a burning candle or lantern. Carbonic 
anhydride is another name for the same gas, it being the 
anhydride of the unstable acid, H 2 C0 3 . It is still popu- 
larly called carbonic acid gas. 

30. Uses of Carbon Dioxide. — Carbon dioxide is used 
extensively in making "soda water." It is confined in 



CARBON AND A FEW COMPOUNDS 145 

strong cylinders under great pressure, and allowed to 
flow into cold water in strong tanks also under pressure. 
The water is thus thoroughly charged. When the stop- 
cock is turned and the water flows into the glass, the 
pressure being removed, the carbon dioxide rapidly bubbles 
out. It is this gas which gives the sharp biting taste to 
soda water and also to the water of many mineral springs. 
It is the same gas that causes the effervescence in beer and 
the sparkling appearance of some wines. 

31. In some of our cities carbon dioxide is now being put 
upon the market in small oval-shaped steel vessels into which 
the gas is forced under great pressure. When ready for 
use, a valve is opened, the gas rushes into a glass of water 
flavored and sweetened, and the soda water is ready. 
These sparklets, as the steel vessels are called, are very 
small, and a large number may be carried without great 
inconvenience. In Germany the same article is sold 
under the name of Sodors. Certain fire extinguishers 
owe their value to the large quantities of carbon dioxide 
contained ; and instances are on record in which fires in 
coal mines which have defied all other means have been 
extinguished by passing in carbon dioxide. 

32. Though this gas cannot be inhaled in any consider- 
able quantities, it is not poisonous, but like water causes 
death by shutting out the oxygen. Hence a person might 
drown in a well or vat of carbon dioxide just as readily as 
in one of water. To plant life, however, the gas is in- 
dispensable ; it is inhaled by plants as oxygen is by 
animals, and in the presence of light the life forces of the 
plant are sufficient to decompose the compound into its 

\ constituents. The carbon is stored up in the woody 
structure of the tree or plant, and the oxj-gen is given off 
again to the air. Thus a considerable portion of the 



146 MODERN CHEMISTRY 

carbon in all our forests and coal beds was once in the 
form of gaseous carbon dioxide in the atmosphere. 

The Hydrocarbons 

33. Definition. — By this term we mean those com- 
pounds consisting of carbon and hydrogen, of which there 
are many. The most important are the three follow- 
ing : — 



Marsh Gas 


CH 4 


Olefiant Gas . 


. . C 2 H 4 


Acetylene 


C 2 H 2 



34. Marsh Gas. — This is also known as methane, and 
by miners as fire damp. It is often found in coal mines as 
the result of the decomposition of organic matter, and in 
swamps from the same source. By stirring the leaves and 
similar matter that collect upon the bottoms of ponds, 
bubbles of gas, consisting largely of marsh gas, are seen 
to rise. It is always produced* in the destructive dis- 
tillation of any organic matter, such as the preparation of 
charcoal in retorts or that of illuminating gas from coal. 

35. Characteristics of Marsh Gas. — Marsh gas is a color- 
less, odorless gas, the lightest of all except hydrogen, 
having a specific gravity compared with air of less than 
0.6. It is highly inflammable, burning with a pale blue 
flame, and with air or oxygen forms a dangerous explosive 
mixture. It is by this gas that most explosions in coal 
mines are caused, and on this account it is called fire damp, 
the word damp with miners being a generic term meaning 
gas. Marsh gas is somewhat soluble in water, and is 
neutral to test paper, affecting neither red nor blue. It is 
an important constituent of ordinary coal gas, and when 
burned produces much heat. 



CABBOX AND A FEW COMPOUNDS 1±1 

36. Protection against Fire Damp. — If you hold a wire 
screen over the flame of a Bunsen burner, you will see 
that the flame does not pass through it, although if you 
bring a lighted match above the screen, you will find 
there a combustible gas. This is because the wire cloth, 
being a good conductor of heat, withdraws it from the 
burning gas and so lowers the temperature that what 
has passed through no longer burns. Now hold the screen 
in the flame until it becomes red hot ; the gas above will 
be ignited and continue to burn. 

An observation of these facts led Sir Humphry Davy 
to design the " safety lamp ? ' which now bears his name. 
It is little more than an ordinary miner's lamp sur- 
rounded by a wire screen. If the miner enters a chamber 
filled with fire damp, though the gas may burn on the 
inside of the screen, there is no danger unless he remains 
until the wire becomes hot enough to ignite the gas 
outside. 

37. Olefiant Gas, C 2 H 4 . — This also is a constituent of 
common illuminating gas, and is formed in the destructive 
distillation of wood and coal. It may be prepared by 
heating ethyl alcohol with sulphuric acid. The latter has 

I strong affinity for water and has the power of abstracting 
it from the alcohol, C 2 H 5 OH or C 2 H 4 , H 2 0, thus : — 

C 2 H 5 OH + H 2 S0 4 = C 2 H 4 + H 2 S0 4 , H 2 0. 

38. Characteristics of Olefiant Gas. — Ethane, as this 
gas is also called, is of about the same density as air, is 
colorless, has a faint odor, and burns with a yellowish 
white light, such as is seen in the ordinary gas jet. It 
is somewhat explosive when mixed with air or oxygen ; 
at -40 atmospheres' pressure it is reduced to a liquid. 



148 MODERN CHEMISTRY 

Acetylene, C 2 H 2 

39. How prepared. — This gas is formed in small quan- 
tities together with other hydrocarbons in the distillation 
of wood and coal. It is prepared now in large quantities 
by treating calcium carbide, CaC 2 , with water, as follows : — 

CaC 2 + H 2 = CaO + C 2 H 2 . 

The lime, CaO, thus formed immediately reacts with an- 
other molecule of water, forming slaked lime, or calcium 
hydroxide, Ca(OH) 2 , thus : — 

CaO + H 2 = Ca(OH) 2 . 

The final reaction then would be indicated by — 

CaC 2 + 2 H 2 = C 2 H 2 + Ca(OH) 2 . 

40. Calcium Carbide. — Calcium carbide is a dark gray 
solid, more or less crystalline in appearance, always giving 
off the odor of acetylene, owing to its decomposition by 
the moisture in the air. In America the greater portion 
of the commercial supply comes from Niagara Falls, where 
it is prepared by fusing at intense heat in electrical fur- 
naces pure lime intimately mixed with charcoal finely 
pulverized. When taken from the furnace, it is packed 
in metallic drums, sealed air-tight, and is then ready for 
shipment. 

Experiment 96. — Into a test-tube put a small lump of calcium 
carbide, cover with water, and quickly insert a cork with delivery tube 
and jet attached. Notice the violent chemical action and the odor of 
the gas. Light the jet and notice with what kind of a flame it burns. 

41. Another Method. — Sometimes this method is varied 
slightly by using a flask fitted with a two-hole cork. 
Through one hole passes the delivery tube, through the 



CARBON AND A FEW COMPOUNDS 



149 



other a funnel with a stop-cock. In this way the flow of 
water can be regulated and the rapid evolution of gas 
prevented. Precaution must be taken in this case not to 
light the jet too soon, as acetylene mixed with air is dan- 
gerously explosive. 

42. Acetylene Generators. — This illustrates one class 
of acetylene generators now offered upon the market, in 
which the water is allowed to drip on the carbide. The 
objection to this is that with the small supply of water 
the carbide becomes so warm as to bring about a partial 
decomposition of the acetylene into other undesirable 
hydrocarbons. 

Experiment 97. — For class-room work an excellent generator 
may be prepared thus : procure a tin can, holding a quart or two, 
and having a screw top. (A can in which some 
varieties of coffee are sold will do.) To the 
inside of the screw top solder a hook. Upon this 
suspend a small bucket or basket made from a 
tin can, and having a wire-cloth or perforated 
bottom. Cut out the bottom of the larger can as 
shown in the cross-sectional view b of the accom- 
panying figure ; then solder two strong bent 
wires, W, W, upon opposite sides of it. 

Xow obtain another can just large enough to 
allow the first to move up and down easily within 
it. Melt or cut out the top ; then cut down two 
flaps about three-fourths of an inch deep, and 
bend them to a horizontal position, as at F. 
Through each flap punch a hole large enough to 
receive the bent guide-wires soldered on to the 
other can. Near the bottom cut a round hole 
and insert a rubber cork, through which passes 
a bent delivery tube extending up nearly to the 
top of the can. 

When ready for work fill this can nearly full 
of water, put some carbide into the basket, sus- 
pend it upon the hook and then lower the first 




W 



150 



MODERN CHEMISTRY 



can into position, the guide-wires passing through the openings in the 
flaps. The screw top enables one to refill the basket without remov- 
ing the entire cylinder. As soon as the carbide touches the water, 
acetylene will begin to form, and, mixed with air, will flow from the 
delivery tube T. 

This generator, which illustrates another class now upon the mar- 
ket, is automatic. In case the delivery tube becomes clogged, the 
increasing pressure of the gas will lift the inner cylinder, and with it 
the basket of carbide, from the water. Or, if the guide-wires become 
caught and prevent this, the pressure on the water will cause it to 
flow out over the flaps. In either case the rapid evolution of gas will 
soon cease. 



,0 



CAUTION. — Before beginning the generation of acetylene 
be sure no lights are in close proximity, and allow the first 
gas generated to escape. It contains too much air for 
good results and is too dangerous. With these precau- 
tions the gas may be used direct from the generator, or 
first passed into an ordinary gasometer, which 
any tinner can make cheaply. 

43. Acetylene Burners. — But to secure steadi- 
ness of flow and safety, it is always better to 
pass the gas through an acetylene burner or 
tip, which differs from the tip of an ordinary 
gas jet only in that instead of a slit there are 
two very small openings drilled, oblique to each 
other. See Fig. 40 ; a is a cross-sectional and 
b the top view. These tips are very cheap, 
and safe because the openings for the exit of 
gas are so small that the flame cannot pass 
back into the generator ; c shows another form of tip 
frequently used.. The two openings compel the issuing 
jets of gas to strike each other obliquely, as in a. 

44. Characteristics of Acetylene. — Acetylene is a color- 
less gas, of an ethereal odor when perfectly pure, but as 




Fig. 40. 



CARBON AND A FEW COMPOUNDS 151 

ordinarly obtained it is very offensive to the smell. It 
is soluble, volume for volume, in water and very explo- 
sive when mixed with oxygen or air. An ordinary jet 
of acetylene burns with a yellowish flame, and owing to 
the large proportion of carbon, — over 92 per cent, — it 
gives off considerable soot. With a burner like the one 
described above it furnishes an intensely white light, rival- 
ing the calcium or Drummond light in brilliancy ; so that 
it is now frequently used for projecting lantern slides 
upon screens and for bicycle lamps. 

45. Intense Heat. — Fine iron wires held in the flame 
are quickly consumed, throwing off sparks as if burning 
in oxygen. 

When used in a blast lamp instead of common gas, acety- 
lene burns with a bluish-white flame. The intensity of 
this is sufficient to melt copper wires readily, and ordinary 
platinum wires in two or three minutes ; furthermore, it 
will even soften porcelain. Iron wires a sixteenth of an 
inch in diameter are quickly fused and burn with a most 
brilliant shower of sparks, especially when a molten 
globule of iron upon the end of the wire is suddenly oxi- 
dized, and being thrown out into the air breaks into a 
> shower of stars. Watch-springs and knife-blades may be 
as easily burned away. In a darkened room the display 
is very beautiful. 

46. Blowpipe for Experiments. 
— The blowpipe best suited to 
this work may be made by almost 
any student. See Fig. 41. The 
outer part, B, is the ordinary black 

; japanned blowpipe, costing only 
a few cents. A hole is cut 
through at the point JE, for the insertion of an inner 




152 



MODERN CHEMISTRY 



tube, which may be made by carefully straightening an 
ordinary eight-inch brass blowpipe. Solder this firmly 
in place, plug the mouth end of the outer pipe with a piece 
of brass through which a small hole has been drilled, and 
the acetylene blowpipe is complete. Connect the inner 
tube with the foot-bellows furnishing the air, and the 
outer tube with the acetylene tank, through the acetylene 
tip. Regulate by means of a stop-cock the flow of gas, 
so that when in operation the acetylene is completely 
burned, with the flame almost entirely blue. 

47. From these experiments it will be seen that the heat 
of this flame is intense, reaching probably 2000° C. 

Experiment 98. — To show the explosiveness of acetylene. In the 
center of the bottom of a pound baking-powder can punch a small 
hole. Place the can, bottom upward, for a minute or so over a tube 
delivering acetylene, then set upon the table in the same position. 
Bring a flame to the touch-hole, when, if the proportions are suitable, 
a violent explosion will ensue, and the can will be thrown several feet 
into the air. If too much acetylene has been introduced, it may burn 
quietly a moment at the opening, until, as more air enters at the 
bottom to take the place of the gas burned, an explosive mixture is 
formed and a report follows. 

Illuminating Gases 



48. One of the most important of these has just been 
considered. It is new as an illuminant, and some problems 
in connection with it have not been entirely solved, but it 
is already being extensively applied in many of the smaller 
towns where no gas plant exists, for railway lighting, bi- 
cycle lamps, etc. . The fact that thus far no appliance has 
been invented for using it in cooking, for the reason that 
the excess of carbon covers the utensils with a deposit of 
soot, has prevented a much more extensive use. 



CARBON AND A FEW COMPOUNDS 



153 



Other Illumixaxts 

49. Besides acetylene, ordinary or coal gas, " water " 
gas, and Pintsch gas deserve notice. 

50. Coal Gas. — This is obtained by the destructive 
distillation of coal in iron retorts. The following diagram 
illustrates the essential features of a gas plant. 




Fig. 42. A Gas Plant. 



51. Preparation. — Soft coal is shoveled into the retort, 
beneath which is the furnace, F. When the retort is 
filled, the door is luted on air-tight. The heat from 
the furnace drives out the gaseous products from the 
coal in the retorts, and they are carried up to the hy- 
draulic main, H. From here the gas is forced by means 
of pumps, not shown in the diagram, through the con- 
densers, a series of pipes several hundred feet in length, 
where it is cooled and the tar condensed. This by-product 
is drawn off by pipes, P, to the tar- well, T, from which it 
is pumped into barrels. 

52. From the condensers the gas goes through the 
scrubber, a large cylindrical tank filled with coke or 
lattice work, over which water slowly trickles, The 



154 MODERN CHEMISTRY 

partition through the center causes the gas to flow down 
one side and up the other ; the coke breaks up the gas 
into bubbles, so as to secure a thorough washing. Here 
the ammonia is mostly removed, and the impure aqua 
ammonia thus obtained is drawn off at intervals, neutral- 
ized with acids, and treated with lime for the preparation 
of the ammonia of commerce, as already described. 

53. The gas next passes through the lime purifiers, a 
number of low cylindrical tanks, containing lime spread 
upon horizontal shelves. The lime dries the gas and at 
the same time removes the sulphureted hydrogen and 
the carbon dioxide. In some works, ferric oxide, Fe 2 3 , 
is used for the same purpose. From the purifier the gas 
passes to the gas-holder, a very large tank, where it is 
stored for use. 

On the inside of the retorts, as previously stated, there 
gradually collects a fine, hard deposit, known as gas car- 
bon, which is now a very useful by-product. 

54. Water Gas. — This gas receives its name from the 
fact that steam is used in one part of the process of manu- 
facture. From the boilers steam is passed into chambers, 
or pipes, containing charcoal or coke, heated red hot. 
Here the vapor and coke react upon each other, the former 
being decomposed, thus : — 

C + H 2 = CO + H 2 . 

Two gases, carbon monoxide and hydrogen, mixed to- 
gether, are thus obtained. Both are combustible, and in 
burning produce great heat, but neither gives any light. 
This mixture, therefore, is next allowed to pass into re- 
torts, kept at a high temperature, into which kerosene, or 
some similar oil, is sprayed. The heat vaporizes and de- 
composes the oil into hydrocarbons that do not liquefy 



CARBON AND A FEW COMPOUNDS 155 

again upon cooling, and which burn with a luminous 
flame. This last step is called " carbureting," and by 
it a gas is obtained not very different in composition from 
coal gas. 

55. Pintsch Gas. — This is the gas so frequently used 
for lighting street cars and railway coaches. It received 
its name from its inventor, who sought to improve the old 
and very unsatisfactory method of lighting coaches in 
England by means of candles. The essential features 
of manufacture are similar to those of the coal gas 
plant. Naphtha is sprayed into retorts heated suffi- 
ciently to decompose the vaporized oil into other hydro- 

• carbons. These are then passed through an improved 
form of condenser, a washer, and lime purifiers into the 
gasometer. 

56. Next the gas is drawn through a cylinder known as 
the " freezer," or " dryer." Here, owing to the action of 
the pumps, it expands, and being cooled thereby, loses all its 
moisture. The same pumps force the gas into large tanks 
called " accumulators," from which it is drawn off into 
smaller tanks for shipment from place to place, or directly 
into the storage cylinders, so frequently seen under railway 

I coaches. This light possesses not only the advantages of 
intensity and whiteness, which coal gas, as ordinarily 
burned, lacks, but unlike ordinary gas, its illuminating 
power is only slightly decreased by strong pressure such 
as is necessary for transportation in storage cylinders. 

57. Natural Gas. — Natural gas is formed by the de- 
composition of organic matter, and the main constituents 
are about the same as those of the other mixed gases used 
for illumination. 

58. Composition of Illuminating Gases. — With the ex- 
ception of acetylene, the illuminating gases noticed are all 



156 MOBEBN CHEMISTBY 

mixtures. The most important constituents of coal and 
" water " gas are given below : — 

Coal Gas 

Hydrogen H, about 46 per cent. 

Marsh Gas .... CH 4 " 38 

Olefiant Gas ... . C 2 H 4 , " 2 

Carbon Monoxide . . CO, " 11 " 

Small amounts of higher hydrocarbons, and such impuri- 
ties as hydrogen sulphide, ammonia, and carbon dioxide. 

Water Gas * 

Marsh Gas ..... CH 4 

Carbon Monoxide . . . CO 
Hydrogen H 

Small amounts of higher hydrocarbons. 



SUMMARY OF CHAPTER 

Classification of free forms of carbon. 
Description, preparation, and uses of. 

a. Diamonds. 

b. Graphite. 

c. Coals. 

Origin of natural coal. 
Varieties of and differences. 
Petroleum and products from it. 

d. Charcoal. 

e. Coke. 

/. Gas carbon. 
g. Lampblack. 

* Water gas contains a larger proportion of carbon monoxide than 
ordinary coal gas; otherwise the two are not very different, 



CARBON AND A FEW COMPOUNDS 157 

Reducing power of carbon. 
Meaning of term. 
Experiment. 
Absorbing power. 

For various substances. 
Practical applications of this power. 
Compounds of carbon. 

The oxides — Names and formulae. 
Preparation of CO. 
Characteristics. 
Sources of C0 2 in the air. 
Laboratory method of preparing. 
Characteristics of C0 2 . 
Experiments to illustrate same. 
Practical uses of C0 2 . 

Soda water, sparklets, sodors, etc. 
Hydrocarbons — Meaning of term. 

Three important hydrocarbons. 
Marsh gas — Sources of, in the air. 
Characteristics. 
Protection against explosions. 
Olefiant gas — Where found. 
Characteristics of. 
Compare with marsh gas. 
Value of each in coal gas. 
Acetylene — How prepared for use. 
Manufacture of carbide. 
Description of acetylene generators. 
Description of acetylene tips. 
Characteristics of acetylene. 

Experiments to show its lighting, heating, and 
explosive properties. 
Other illuminating gases. 
Coal gas. 

Method of preparing — Apparatus. 

Plans for purifying. 

Different forms of gas-burners. 

Valuable by-products — How secured — Use. 



158 MODERN CHEMISTRY 

Water gas. 

How prepared. 

Characteristics of. 

Comparison with coal gas in composition. 
Pintsch gas. 

Method of preparing and purifying. 

Used where. 

Comparison with coal gas. 



CHAPTER XII 

FUNDAMENTAL LAWS OF CHEMISTRY 

1. Quantitative Work. — It may have seemed to the 
student that the quantity of a reagent used in any experi- 
ment makes little difference. While definite amounts are 
usually specified, care is not often taken to use exactly 
that quantity. Generally the result will be the same, but 
if more than the necessary amount of a substance is used, 
the excess remains and is simply wasted. This fact is 
usually stated in what is known as — 

2. The Law of Definite Proportions. — Briefly, it is this: 
Two or more elements, in uniting to form a compound, always 
do so in the same proportion by weight. This has been 
illustrated somewhat in the earlier part of the book in 
discussing compound bodies. It is a very important law, 
and upon it much of the science of chemistry depends. 
To illustrate it more fully the student should make the 
following experiments, using the utmost care to insure 
accuracy. Let him not draw his conclusions beforehand 
and then endeavor to make his results conform to these. 

Experiment 99. — Fill two burettes, one with a solution of caustic 
soda and the other with dilute hydrochloric acid, and support them 
upon a stand. Carefully take the reading of each, using the lowest 



FUNDAMENTAL LAWS OF CHEMISTRY 



159 



-37 



—39 



Fig. 43. 



part of the meniscus in doing this, as shown in the figure. Here 
the lowest part of the curve coincides with 38.4, and this would be 
the reading. 

Now find the weight, as accurately as possible, of a small 
evaporating dish. Much time can be saved here if each 
student will provide himself with a small pasteboard box 
and cover, such as blank labels are packed in. Put the box 
and cover upon the scale pan opposite to the dish, and pour 
in fine shot or sand until it is exactly counterpoised. This 
represents the weight of the dish. Put the box with its 
contents away where it will be safe from accident. 

Xow, from the caustic soda burette allow 10 cc. to flow 
into the evaporating dish, and add one drop of phenolphthal- 
ein solution, or, if more convenient, enough litmus solution 
to give a decided blue color. From the other burette, with 
constant stirring, let the acid flow in slowly until the color given 
by the phthalein barely disappears, or until the blue litmus just shows 
pink. Take the reading of the acid burette, and by subtracting 
,the previous reading determine how much hydrochloric has been 
used. The change in the color noted above indicates that sufficient 
acid has been added to neutralize the alkali and form therewith 
a salt, 

Xow place the evaporating dish upon a ring-stand, or better upon 
a sand-bath, and evaporate slowly to dryness. Do not let the liquid 
boil, as some will be lost by spurting out, and be careful toward the 
close to withdraw the heat before the solution is entirely dry, lest the 
( dish become so warm as to decompose some of the salt. If the heat 
of the dish does not complete the evaporation, warm it very gently for 
another moment. When perfectly dry let the dish cool, and weigh it. 
In doing this put the small box and shot upon the opposite pan as be- 
fore, then whatever weights are necessary to add will represent the 
weight of the salt obtained. If the shot are not used, subtract the first 
weight from the second. Tabulate results as below : — 

Caustic soda used . . 10.0 cc. 
HC1 used 6.1 cc. 



Wi 



t. of dish + salt 
Wt of dish . . 

Wt of salt . . . 



17.103 
15.217 

1.886 



Experiment 100. — Purpose, a continuation of the preceding. 
Repeat the preceding experiment, using the same amount of caustic 



160 MODERN CHEMISTRY 

soda, but twice as much acid. The litmus or phthalein need not be 
added. Use the same precautions as before. Tabulate results. 

Wt. of dish + salt . . Caustic soda used . . 10.0 cc. 

Wt. of dish .... 15.217 HCl used ..... 12.8 cc. 
Salt 

Experiment 101. — Purpose, same as above. Repeat, using 5 cc. 
of caustic soda solution, a few drops of litmus or one of phthalein, and 
then enough hydrochloric to neutralize, as in Experiment 98. Cool 
and weigh as before. 

3. Comparison of Results. — Comparing the results ob- 
tained, we may* formulate them as below : — 

10.0 cc. Salt (NaCl) obtained . . — 
6.4 cc. 



rjJLy. vv. 


HCl " . 


Exp. 100. 


NaOH used . 




HCl " . 


Exp. 101. 


NaOH used . 




HCl " . 



10.0 cc. NaCI obtained 
12.8 cc. 
5.0 cc. NaCI obtained 



4. What evidence in the above experiments do you find 
in proof of the law of definite proportions ? Is there any 
agreement between the first and second of the above? 
Between the second and third ? Why ? 

Experiment 102. — Further proof of the law. Carefully weigh an 
evaporating dish or find its equivalent in shot as before, then add a half- 
gram weight to the pan on which the shot is, and put into the evaporat- 
ing dish sodium carbonate crystals to balance. Add a few centimeters 
of pure water to the carbonate, and then add dilute hydrochloric acid, a 
little at a time. Keep the dish covered with a sheet of glass or watch 
crystal so as not to lose any by its spattering out. In this way cau- 
tiously add the acid until the carbonate is all dissolved, or until it no 
longer effervesces. Now rinse off the cover-glass into the evaporating 
dish, and evaporate to dryness with the same precautions used before. 
Cool, weigh, and determine the amount of salt obtained. 

Sod. Carb. : Na 2 C0 3 + dish . . — NaCl + dish . . 
Wt. of dish ... — Dish 



Na 2 C0 3 — NaCl 



FUNDAMENTAL LAWS OF CHEMISTRY 161 

Experiment 103. — Same as preceding. Pursue the same method 
as above, using 1 g. of the carbonate instead of a half gram. 

Used: Na 2 C0 3 + dish . . — Obtained: XaCl + dish . . — 
Wt. of dish .... — Dish — 



Xa 2 C0 3 — NaCl — 

Experiment 101. — Purpose, same as before. Repeat the preced- 
ing, using this time 1^ g. of sodium carbonate crystals. Results : — 

Used : Xa 2 C0 3 + dish . . Obtained : XaCl + dish . . — 

Dish — Dish — 



Xa C0 3 1.500 XaCl 



5. Summary. — In each of the last three experiments 
find the ratio existing between the carbonate used and the 
salt obtained. 



1. Na 2 C0 3 :NaCl 

2. Na 2 C0 3 TNaCl 

3. Na 9 C0 Q :NaCl 



:i '2 vv 3 •• • z — 



l:y 



Is there any uniformity in the value of these ratios ? 
Do your results afford further evidence of the law of 
definite proportions ? If so, in what way ? 

6. The Law of Multiple Proportions. — We have learned 
that when two or more elements unite to form a compound 
they do so in a constant ratio. We have seen, however, 
that the same two or three elements may unite to form 
several compounds, and at first this may seem contrary to 
the statement of the preceding law. It is a modification, 
but not a contradiction. If a new and different compound 
is formed when other proportions are used, in this the 
quantity of the elements that enter into combination is 
always some multiple of the lowest. An illustration will 
make this plain. Thus, we are familiar with the series of 
nitrogen oxides : — 



162 



MODERN CHEMISTRY 



Nitrogen Monoxide . 


. N 2 


N 2 


:0 : 


:28 


16 


" Dioxide . . 


• N 2 2 


N 2 


2 : 


:28 


32 


" Trioxide . 


• N 2 3 


N 2 


3 : 


:28 


48 


" Tetroxide . 


• N 2 4 


N 2 


<V 


:28 


64 


" Pentoxide 


• N 2 5 


N 2 : 


<V 


:28 


80 



It is seen that while the weight of the nitrogen entering 
into combination remains constant, the oxygen is in the 
ratio of 2, 3, 4, and 5 times what it is in the lowest of the 
series. 

7. This law may be proved experimentally by estimat- 
ing the amount of oxygen that a given weight of potas- 
sium chlorate, KC10 3 , will yield, by the method previously 
suggested. Then, determine the amount in potassium 
perchlorate, KC10 4 . In these two compounds we should 
find the ratio to agree with that expressed in the formulae, 
that is, 3 and 4 times what would be contained in a mole- 
cule like mercuric oxide, HgO. 




Fig. 44. 



Experiment 105. — To prove the law, the work may be conven- 
iently done as shown in the accompanying figure. Instead of the 
flask, 0, a hard-glass test-tube may be used. Put into the flask about 
1 g. of manganese dioxide, Mn0 2 , and weigh carefully the flask and 



FUNDAMENTAL LAWS OF CHEMISTRY 168 

contents. Then add to it about \\ g. of potassium chlorate, KC10 3 , 
and weigh accurately. The difference between the two weights will 
be the exact amount of the chlorate used. A is a two or three liter 
bottle with narrow neck, and fitted with a two-hole rubber cork. The 
delivery tube, d, just passes through the corks of and A ; the con- 
necting tube, e, reaches nearly to the bottom of A. 

The bottle, B, is similar to A, but is not provided w T ith a cork, and 
the tube, e, need not extend far into the bottle. When the connec- 
tions are made air-tight, gradually heat the flask, 0, thus driving the 
oxygen over into A, and the water from A into B. Continue the 
operation until no more gas is given off, then allow the whole to cool 
to the temperature of the room. Now measure the water in the bottle, 
B\ this will give the volume of oxygen at the temperature of the room 
and the prevailing barometric pressure. According to methods already 
given on page 96, reduce this volume to what it would be under stand- 
ard conditions. Knowing the weight of a liter of oxygen, 1.43 g., find 
the weight of the determined volume. As a check, weigh also the 
cooled flask, 0, and determine its loss ; this should be practically the 
same as the weight found by calculation. 

Next, arrange the apparatus as at the beginning. Into the hard- 
glass tube put about 1.4 g. of potassium perchlorate, KC10 4 , and, after 
making connections, heat strongly as before until no more gas is pro- 
duced. Cool and weigh the flask or tube ; the loss will represent the 
oxygen, which may be checked up by determining the volume of 
the gas given off as before and reducing to standard conditions. 
Let the student now compare results. The two reactions are as 
follows : — 

KC10 3 + heat = KC1 + 3 0, 

and KC10 4 + heat = KC1 + 4 0. 

From other experiments we know that the oxygen is entirely removed 
and that potassium chloride, KC1, remains. Then we should have the 
proportion 

Mol. wt. KCIO3 : wt. of O in 1 mol. KC10 3 : : 1.25 g. : m g., 

, in which m = no. grams found by experiment above. Substituting, 

122.5 : x : : 1.25 g. : m g., 

m x 122.5 
x = . 

* 1.25 



164 MODERN CHEMISTRY 

Then, as 16 is the weight of one atom of oxygen, there would be as 
many atoms of oxygen in the molecule as 16 is contained times in x\ 
the result should agree very closely with the assumed number. 
In the same way, 

Mol. wt. KC10 4 : wt. of O in 1 mol. KC10 4 : : 1.40 g. : n g., 

in which n = no. grams found in second instance above. Or, 

138.5 : y : : 1.40 g. : n g., 

= n x 138.5 
V ~ 1.40 
How does the value of y agree with the known value? 

8. Combining Weights. —We have previously learned 
that when elements or compounds react with each other 
in the formation of new substances, they always do so in 
a fixed or definite proportion. We have seen also that 
when several compounds are formed from the same two 
elements, there is one smallest quantity of which all the 
others are multiples. This smallest amount in the case of 
the nitrogen oxides was found to be 16 for the oxygen, 
and all the others were multiples of this. Therefore, 16 
is regarded as the atomic weight of oxygen, and in all 
chemical reactions into which it enters, this, or some 
multiple of it, is its combining weight. 

Experiment 106. — To find the combining weight of copper. Put 
into a beaker 2 \ g. of clean, bright copper, accurately weighed, and 
dissolve slowly in nitric acid somewhat diluted. Use every precaution 
to prevent loss by spurting, just as in other similar work, and when 
the copper is all dissolved, transfer the solution to a weighed evapo- 
rating dish, as small as will conveniently hold the solution ; carefully 
rinse off the cover-glass and the beaker into the evaporating dish, and 
evaporate to dryness. We now have a blue salt, copper nitrate. Be 
sure it is perfectly dry, and then remove the sand-bath, or any other 
protection used for the dish, and gradually increase the heat until all 
particles of the blue salt have been changed to a black compound. A 
dull red heat is generally necessary for this. We now have copper 



FUNDAMENTAL LAWS OF CHEMISTRY 165 

oxide, CuO. Cool and weigh carefully. Determine how much oxygen 
has combined with the copper by subtracting the amount of copper 
used from the weight of oxide obtained. 

Dish + CuO . — m + x ; m = wt. of dish ; x, of CuO. 
Dish + Cu . m + n; n = wt. of Cu = 2\ g. 

O m + x — (m + n) — y. 

Numerous experiments have shown that the combining weight of 
oxygen is 16 ; using this as a basis, we can determine w T hat it is for 
copper : — 

Wt. of O found : wt. of Cu used : : comb. wt. of O : comb. wt. of Cu. 

or, Wt. of O : 2| g. Cu : : 16 : z. 

From this proportion the combining weight of copper should be found 
to be approximately what is given in the. table on page 9. The sources 
of error are liable to make the difference comparatively great, but the 
result should not vary too much. 

Experiment 107. — Purpose, same as above. Use 3 or 4 g. of 
finely powdered copper nitrate which has not been exposed to the air 
any length of time. Be sure the exact weight is known, then heat 
in a small evaporating dish, or better, in a porcelain crucible, cautiously 
at first; when the nitrate is converted into the black oxide as before, 
cool and find the weight. Experiment has shown that 1 g. of crystal- 
lized copper nitrate contains 0.2619 g. of metallic copper. From this 
determine the amount of copper represented by the 3 g. (or 4 g.) of 
nitrate used. 



Wt. of dish + CuO . . . . 






Wt. of dish 












CuO 




Cu 

















Wt. of : wt. of Cu : 


: 16 : x. 


. 



Does this give practically the same combining weight for copper 
that the preceding did ? If the results do not correspond fairly well 
with each other and with the table, the experiments should be repeated. 
Experiment 108. — To find the combining weight of tin. For this 
use granulated tin. If not at hand, procure a quantity of pure tin 



166 MODERN CHEMISTRY 

foil, melt it in an iron ladle, and pour into cool water. Remove from 
the water and dry it, when it will be ready for use. Weigh out care- 
fully 2 g. of the granulated tin, and treat with nitric acid in an evapo- 
rating dish. Take care always to avoid loss by spurting. Evaporate 
slowly to dryness, and then gradually heat the white residue to dull 
redness. 

Wt. of dish + Sn0 2 — 

Wt. of dish — 

Sn0 2 — 

Wt. of Sn — 

2 — 

It has been found by analysis that the amount of oxygen in this 
compound indicates two atoms to the molecule, hence in making our 
calculations that amount must be used. Then we have — 

wt. of O found : wt. of Sn used : : 32 (2 x 16) : x. 

Experiment 109. — Repeat the above experiment, using 2| or 3 g. 
of tin, and make calculations as before. 

How do the results in the two experiments agree? If they do not 
correspond fairly well with the atomic weight given in the table, 
allowing for errors in weighing, the experiment should be repeated. 

9. Such experiments as the above might be endlessly 
multiplied. We have found in these, as has been the 
case in an indefinite number of instances in which chem- 
ists have done the work with the utmost care, that every 
element combines with others in some exact proportion 
by weight, and whether we use much or little of the 
element, in the same compound the ratio never changes. 
This fact is of the utmost importance, for upon it depends 
much of the science of chemistry. It is this that enters 
into the application of chemistry to the arts and manufac- 
tures, and renders its results so sure and unchanging. 

10. Some Application of the Laws of Combination. — 
Knowing that the laws of combination are true, Ave may 
make use of the principles in determining the strength 



FUNDAMENTAL LAWS OF CHEMISTRY 167 

of acid or alkaline solutions. The following work will 
illustrate this. 

Experiment 110. — To determine the strength of any hydrochloric 
acid solution in the laboratory. Put the acid to be tested into a 
burette and take the reading. From this allow 10 cc. to flow into an 
evaporating dish, add a drop or two of litmus or phthalein, and then, 
from another burette, after taking the reading, run in a solution of 
caustic soda until the solution in the evaporating dish is neutralized, 
as in previous work. Evaporate slowly to dryness, cool, and weigh. 
Subtract the weight of the dish to determine the salt obtained. Sup- 
pose this is 0.585 g. Xow we know that caustic soda and hydrochloric 
acid react with each other according to the following equation : — 

NaOH + HC1 = NaCl + H 2 0. 

From this w T e see that one molecule of pure hydrochloric acid yields 
one of sodium chloride, or 36.5 parts by weight of acid give 5S.5 of 
salt, The 0.585 g. of salt would thus correspond to 0.365 g. of acid, 
the amount in 10 cc. of the solution used. Then in a liter, 1000 cc, 
there would be 100 times this amount, or 36.5 g. of pure acid. The 
liter of acid then ought to weigh 1000 g. + 36.5 g. = 1036.5 g. The 
question simply is this : 36.5 g., the amount of pure acid, is what 
per cent of 1036.5 g., the weight of the acid solution? This is found 
to be about 3J per cent. 

Experiment 111. — Repeat the preceding experiment, neutralizing 
10 cc. of the hydrochloric acid with caustic potash. Make your cal- 
I culations from the following equation : — 

KOH + HC1 = KC1 + H 2 0. 

Do your results agree with the preceding as to the per cent strength 
, of acid? 

11. To determine the Amount of Caustic Soda or Potash 
in the Solutions used above. — We know that when an acid 
and an alkali are put together, they neutralize each other 
i to form a salt. If then we know how much acid is con- 
tained in a solution, and measure the amount of the latter 
used, having some means of knowing when the alkali is 



168 MODERN CHEMISTRY 

exactly neutralized, we can easily calculate the amount of 
alkali contained in a given volume of solution. 

Experiment 112. — Suppose we are required to determine the 
number of grams of sodium hydroxide in 1 liter of the solution. We 
know the reaction is 

NaOH + HC1 = NaCl + H 2 0, 

or by weight, 40 + 36.5 = 58.5 + 18. 

That is, 40 g. of caustic soda are necessary to neutralize 36.5 g. of 
hydrochloric acid. Suppose now we have a solution of acid that con- 
tains 3.65 g. of pure hydrochloric acid to the liter, then 

1000 cc. HC1 would neutralize 4.0 g. of NaOH. 

Then, if with 100 cc. of caustic soda solution we used 20 cc. of the 
acid solution, we should have this proportion : — 

1000 cc. HC1 : 4.0 g. NaOH : : 20 cc. HC1 : x g. NaOH ; 

x = .08. 

That is, in 100 cc. of the solution of caustic soda there are .08 g. of 
the solid alkali dissolved ; then in 1 liter there would be 10 times as 
much, or .8 g. 

For such work as this we ordinarily use oxalic acid instead of 
hydrochloric, because it is easier to obtain pure, and forms a more 
constant solution. Its formula is H 2 C 2 4 , 2 H 2 0. With caustic soda 
the reaction is 

2 H 2 0, H 2 C 2 4 + 2 NaOH = Na 2 C 2 4 + 4 H 2 0, 

or by weight, 126 + 80 = 134 + 72. 

That is, 126 g. of oxalic acid will neutralize 80 g. of caustic soda. 
Suppose for work we weigh out 6.3 g. of oxalic acid and dissolve in 
1000 cc. of pure water. This will be our standard solution of acid. 

To find how much caustic soda in 1 liter of solution. Measure out 
accurately into a beaker 50 cc. of the alkali solution, and add one drop 
of phenol phthalein, or about 1 cc. of litmus solution. Next take the 
reading of a burette containing the standard oxalic acid solution, and 
with constant stirring let the acid drop in slowly until, finally, by the 
addition of a single drop the red color of the phenol disappears, or the 



FUNDAMENTAL LAWS OF CHEMISTRY 169 

blue of the litmus changes to red. Again read the burette and deter- 
mine how much acid has been used. Suppose it has been 10 cc. Then 
to calculate, 

1000 cc. acid ■ 4.0 g. NaOH : : 10 cc. : x NaOH ; 

x = .04 g. NaOH, 

the amount in 50 cc. of NaOH solution used. In 1000 cc. there would 
be 20 times as much, or .8 g. 

Problem 1. — Let the teacher make up a solution of caustic potash 
with distilled water, and have the student determine the number of 
grams used to the liter. 

Problem 2. — In the same way let the student determine the 
amount of common salt in a solution by using in the burette a 
solution of silver nitrate containing 17 g. per liter. To determine 
' when sufficient silver nitrate is used, add to the common salt solution 
sufficient potassium chromate solution to give a yellow color. With 
constant stirring run in the silver nitrate until the precipitate that 
forms shows the faintest red tinge. The reaction is 

NaCl + AgN0 3 = AgCl + NaN0 3 , 
or by weight, 58.5 + 170 = 143.5 + 85. 

That is, 170 g. of silver nitrate will precipitate the chlorine in 58.5 g. 
of salt, or if 17 g. of silver nitrate were used to make a liter of the 
solution, then 1000 cc. of silver nitrate would precipitate the chlorine 
in 5.85 g. of salt. Or, 



1000 cc. AgNO s : 5.85 g. NaCl ::mcc, AgN0 3 ■ x g. NaCl, 



in which m is the number of cubic centimeters of silver nitrate solu- 
tion used with the amount of common salt solution taken. If this 
I latter is 20 cc, or -^ of a liter, then 50 m — number cubic centimeters 
AgX0 3 necessary to precipitate the chlorine in 1 liter. 

12. Displacing Power of Metals. — We have seen in pre- 
paring hydrogen that various metals have the power of 

I reacting with certain acids to displace the hydrogen 
contained. Of course this displacing power is in accord- 

1 ance with the valence of the element (see chapter on 



170 



MODERN CHEMISTRY 



Valence), and the following plan maybe used to deter- 
mine it : — 

Experiment 113. — Into a 4-ounce wide-mouth bottle, C, put 1 
gram of finely granulated zinc. Insert a cork fitted with a safety 
tube, A, running to the bottom of the bottle, 
and with the delivery tube, B. Let the deliv- 
ery tube extend over to a pneumatic trough, 
and place the end of the tube under the mouth 
of a 2-liter bottle filled with water. When 
everything is in readiness, by means of the 
funnel-tube, add to the zinc in C about 30 or 
40 cc. of moderately strong hydrochloric acid. 
When the zinc is all dissolved and the appara- 
tus cooled, measure the amount of hydrogen 
collected, by again filling the receiver with 
Fig. 4o. water, noting the amount required. Reduce 

this volume to standard conditions. Suppose, for example, it is found 
that it requires m cc. of water to refill the receiver, and that by reduc- 
ing this volume to standard conditions, we obtain n cc. as a result, 
then as 1 liter of hydrogen weighs .0896 g., n cc. would weigh 




n x .0896 
1000 



= w g. of hydrogen. 



Then we should have the proportion 

to g. of H. : 1 g. of Zn : : x : 65 ; 

that is, the weight of the hydrogen obtained is to the weight of the 
zinc used in displacing it as x is to 65, the atomic weight of zinc. 
This should give for the value of x, approximately, 2. Then as the 
hydrogen atom is the standard, or 1, in this case x represents the 
weight of two atoms of hydrogen. In other words, the zinc atom has 
the power of displacing two atoms of hydrogen. 

Experiment 114.— With apparatus arranged as in the preceding 
experiment, let the student use one gram of aluminum wire cut into 
small pieces. No heat will be necessary if strong hydrochloric acid 
be used, and the chemical action, slow at first, will soon become very 
rapid. Determine as before the volume and weight of the hydrogen 
set free. Then we have 



SULPHUR AND ITS COMPOUNDS 171 

wt. of H obtained : wt. of Al used : : x : 27, atomic wt. of Al, 
and x = ? 

From this what can you say is the displacing power of the aluminum 
itom ? 

Experiment 115. — In exactly the same way try 1 gram of magne- 
sium ribbon, cut into small pieces. Hydrochloric acid somewhat 
diluted had better be used, as the action is very rapid. Make your 
corrections for temperature and pressure, and calculate as before. 
What do you find for the displacing power of the magnesium atom ? 

SUMMARY OF CHAPTER 

Statement of Law of Definite Proportions. 

Experiments illustrating it. 
Law of Multiple Proportions. 

How illustrated. 
Combining weights. 

Method of determining by experiment. 
For copper. 
For tin. 
Practical application. 

Method of determining amount of acid or alkali in a solution. 

Method of determining valence or displacing power of metals. 



CHAPTER XIII 

SULPHUR AND ITS COMPOUNDS 

1. Where found. — Sulphur is an element that has been 
known from very early times. By some of the alchemists 
it, together with mercury, was regarded as forming all of 
the metals. 

It is a native of volcanic regions, and is found in abun- 
dance in Sicily and to some extent in Iceland. There are 
said to be some deep deposits in the Southern States, but 



172 



MODERN CHEMISTRY 



they have not been developed. In the form of compounds 
with the metals, sulphur is found abundantly and very 
widely distributed. Some of the more common of these 
compounds are gypsum,, CaS0 4 , iron pyrite, FeS 2 , galena, 
PbS, and zinc blende, ZnS. In the form of hydrogen sul- 
phide, it is found in many mineral springs and is often 
emitted from volcanoes. 

2. For many years Sicily had a monopoly of the sul- 
phur trade. It occurs there in almost unlimited quan- 
tities, mixed with earthy matter. This mixture may be 
partially purified by a method similar to that employed 
in the preparation of charcoal. Large piles of the crude 
sulphur are heaped up and covered with earth and sod. 







Fig. 46. 



The mass is then ignited and a part of the sulphur in 
burning melts the remainder, which runs out into trenches | 
or vats, leaving the earthy matter behind. 



SULPHUB AND ITS COMPOUNDS 173 

3. For many purposes the sulphur thus obtained needs 
further purification. It is heated and vaporized in retorts, 
the vapors passing over into cool chambers and condensing 
upon the walls in the form known as flowers of sulphur. 
If the operation continues for a length of time, however, 
the walls become heated enough to melt the sulphur that 
forms upon them. It is then allowed to run out into 
molds, in which form it is known as brimstone or roll 
sulphur. (See Fig. 46.) S is a cylinder in which the 
sulphur is melted, F", a retort where it is vaporized, and 
E. the condensing chamber. 

4. New Source of Supply. — From the fact that Sicily 
controlled the sulphur trade, prices rose so high at one 
time that the English manufacturers were obliged to 
resort to some other source of supply. Sulphur was 
used extensively in making sulphuric acid for the manu- 
facture of soda crystals. It was found that by roasting 
iron pyrite, FeS 2 , a compound that had been hitherto alto- 
gether worthless, the sulphur dioxide could be obtained ; 
or if the ore was heated in retorts sealed up to prevent 
access of air, the sulphur was not oxidized, and could 
be condensed. As the pyrite is very abundant, and the 
method of obtaining the sulphur cheap, this at the present 
time furnishes not only about all the sulphur needed in 
making sulphuric acid, but even more, so that the demand 
for Sicilian sulphur has greatly decreased. The reaction 
that takes place when iron pyrite is heated in sealed 
retorts is g ¥q ^ + heat = 2 S + Fe 3 S 4 . 

5. Characteristics of Sulphur. — Sulphur is a yellow, 
brittle solid, a little heavier than water. It is seen in a 
number of forms, of which the flowers and roll sulphur 
have been mentioned. It also occurs crystallized. 



174 MODERN CHEMISTRY 

Experiment 116. — Into a test-tube put about a cubic centimeter 
of carbon disulphide, and add a little sulphur. When the latter has 
dissolved, pour off the clear solution upon a watch crystal, and allow 
it to evaporate slowly to dryness. The sulphur will form in crystals, 
the shape of which may be recognized if the evaporation is slow. If 
necessary, however, examine with a magnifying glass. What form 
have they ? 

Experiment 117. — Fill a small crucible nearly full of sulphur, 
and heat till it is melted. Allow it to cool, and when a crust has 
formed over the surface, break an opening in the top and pour out 
what remains molten. Let it cool a little more and break open the 
mass. What kind of crystals have formed? 

Experiment 118. — Put 4 or 5 g. of sulphur into a test-tube and 
warm. Note how it changes, first melting to form a light yellow- 
colored liquid, then becoming quite thick again and very dark, then 
thin again. At this last stage, pour out the sulphur into cold water 
and note its condition. This is called amorphous sulphur, or some- 
times plastic sulphur. 

6. Sulphur is found in a large variety of crystallized 
forms. The octahedral and the long, needle-like crystals 
have been seen. Upon standing for some time these 
gradually change into other forms, modifications of the 
two. Amorphous sulphur is dark-colored and elastic, 
somewhat like rubber. It is regarded as an allotropic 
form. Sulphur is insoluble in water, hence has no taste ; 
it is also without odor. As we have seen, it is soluble in 
carbon disulphide. It is combustible, burning with a pale 
blue flame, and forming the well-known irritating gas, 
sulphur dioxide. At high temperatures sulphur combines 
readily with most of the metals, forming sulphides. This 
has been shown already in preparing ferrous sulphide by 
heating iron filings mixed with sulphur. Copper turnings 
serve equally well. 

7. Comparison of Ozone with Allotropic Sulphur. — In 
the case of ozone, we have seen that its molecule is differ- 



SULPHUR AND ITS COMPOUNDS 175 

ent from that of the oxygen molecule. The same is be- 
lieved to be true of sulphur and its allotropic form, as well 
as of all other elements which show the same variation. 
We cannot prove this for sulphur, but there are some 
facts which make this theory strongly plausible. Thus, 
if the vapor is weighed at 1000° temperature, it is only 
one-third as dense as when weighed at 500°. 

8. Uses of Sulphur. — Sulphur is used largely in the 
manufacture of gunpowder, the other two constituents 
being charcoal and saltpeter. These are united in about 
the following proportions : — 

Sulphur 12 per cent 

Charcoal 13 " 

Saltpeter 75 " 

" Greek fire," which played so important a part in the 
early centuries, and the composition of which was kept a 
secret for several hundred years, differed very little from 
the gunpowder of the present time. Sulphur is employed 
to some extent in the manufacture of rubber goods, espe- 
cially vulcanite, and considerably in fumigating buildings; 
it is used largely in making sulphuric acid. Because of 
its low kindling-point sulphur has been used very exten- 
sively in the manufacture of matches, but the irritating gas 
produced, and the slowness with which such matches burn, 
have led to the substitution of other substances. 

Compounds of Sulphur 

9. Hydrogen Sulphide, H 2 S. — This gas, known also as 
sulphureted hydrogen, occurs in many mineral springs, 
which give it off abundantly; it is sometimes emitted 
from volcanoes, and is noticed in the decay of eggs and 
other similar substances. 



176 MODERN CHEMISTRY 

10. How prepared. — For laboratory purposes hydrogen 
sulphide is always prepared by treating ferrous sulphide, 
FeS, with sulphuric acid. 

Experiment 119. — Owing to the offensive odor of the gas, it should 
be prepared in very small quantities, and kept from access to the room 
as much as possible. Put into a test-tube a small bit of ferrous sul- 
phide, cover it with water, and add a few drops of strong sulphuric 
acid. Action will begin at once. Notice the odor of the gas. Has 
it any color ? Attach a jet and ignite it. With what kind of a flame 
does it burn? Notice the odor given off by the burning gas. Hold a 
cold beaker over the flame. What do you see depositing upon it? 
What are the two products formed when hydrogen sulphide burns? 
Write the reaction. 

The reaction that takes place when hydrogen sulphide is prepared 
is seen below : — 

FeS + H 2 S0 4 --= FeS0 4 4- H 2 S, 

FeS + 2 HC1 = FeCl 2 + H 2 S. 

Hydrochloric acid may be used instead of sulphuric acid with good 
results. 

11. Characteristics of Sulphureted Hydrogen. — It is a 

colorless gas, having a very disagreeable, nauseating odor ; 
is somewhat poisonous, and should not be inhaled. It is 
inflammable, burning with a bluish flame, is a little 
heavier than air, and somewhat soluble in water. It has 
the power of forming precipitates with solutions of many 
metallic salts. 

Experiment 120. — Into a test-tube put a little of a mercuric 
chloride solution, into another a solution of antimony tartrate, into 
a third arsenic trioxide dissolved in hydrochloric acid and water. 
Attach a delivery tube to a hydrogen sulphide generator, and pass the I 
gas through each of the solutions. Notice the color of the precipitates 
obtained. Lead salts are very sensitive to the action of hydrogen 
sulphide, and are used in testing for its presence. 

12. Use of Hydrogen Sulphide. — Mineral waters con- 
taining this gas in solution are supposed to be beneficial 



SULPHUR AND ITS COMPOUNDS 177 

to health. With this exception, about the only use for 
hydrogen sulphide is in the laboratory, as a reagent, espe- 
cially in making analyses of unknown solutions. Many 
of the metals in the form of salts are converted by hydro- 
gen sulphide into insoluble sulphides. Such metals, there- 
fore, when treated with the gas, may be separated from 
Dthers which are not so precipitated. 

Experiment 121. — The above statements will be made plain by 
:his experiment. Into a beaker put a few cubic centimeters of a solu- 
:ion of mercuric nitrate and as much of zinc sulphate ; add a few drops 
Df hydrochloric acid, and pass through it a current of sulphureted 
rydrogen until the odor of the gas is still perceptible after shaking 
;he solution. Then filter out the black precipitate and test the clear 
iltrate for zinc with ammonia, as you have done in Chapter VII, Sec- 
don 12. Have you succeeded in separating the two metals ? 

13. Oxides of Sulphur. — There are two of these com- 
pounds, the dioxide and the trioxide, S0 2 and S0 3 . It 
is only the first that is of special importance or interest 
:o us. 

14. Sulphur Dioxide, S0 2 . — This is also known as sul- 
phurous anhydride, because by passing it into water sul- 
ohurous acid is formed. It is the familiar, irritating gas 
always produced when sulphur is burned in the air. 

15. How prepared. — For laboratory purposes sulphur 
iioxide is prepared by treating copper turnings with 
strong sulphuric acid. The reaction is usually indicated 
as follows : — 

Cu + 2 H 2 S0 4 = CuS0 4 + 2 H 2 + S0 2 . 

If this is compared with the reaction of zinc and sulphuric 
acid upon each other, it will be seen to be very different. 
Zinc is acted upon by cold, dilute acid, while copper re- 



178 MODERN CHEMISTRY 

quires the acid hot and concentrated. It is probable that, 
as with zinc, hydrogen is first formed, thus : — 

Cu + H 2 S0 4 = H 2 + CuS0 4 , 

and that this nascent hydrogen immediately attacks an- 
other molecule of sulphuric acid, decomposing it, thus : — 

H 2 S0 4 + H 2 = 2 H 2 + S0 2 . 

Putting these two reactions together, we have the one 
given above. 

Experiment 122. — Put into a test-tube a few copper turnings and 
nearly cover with strong sulphuric acid. Heat moderately until the 
fumes begin to come off, and collect two or three bottles of the gas as 
you have carbon dioxide, by downward displacement. What is the 
odor of the gas ? Test it to learn whether it will support combustion 
or will burn. What can you say of its density ? Try its effect upon 
moistened red and blue litmus paper; state results. Pour into one 
bottle of the gas a few cubic centimeters of litmus, cochineal, or some 
other vegetable solution, and shake it. What happens? Suspend in 
another bottle some colored paper, or silk or straw goods, moistened, 
and allow to remain some time. State results. 

Invert another bottle or test-tube filled with sulphur dioxide, over a 
small evaporating dish of water. Does the water rise in the tube? 
Why? Test the water with blue litmus paper; what effects? What 
has been formed with the water ? 

16. Characteristics of Sulphur Dioxide. — It is a very 
irritating, colorless gas, considerably heavier than air. It 
is soluble in water, forming an acid solution, which, how- 
ever, is very unstable. It will neither burn nor support! 
combustion, though magnesium ribbon will burn in it with 
difficulty as it does in carbon dioxide. It is. readily lique- 
fied by passing the gas through a spiral tube, surrounded 
by ice and salt. In the liquid condition it is limpid, trans- 
parent, and very slightly yellow in color. 



SULPHUR AND ITS COMPOUNDS 179 

17. Sulphur dioxide is a great reducing agent, like 
charcoal, but more active. That is, it has the power of 
abstracting oxygen from other substances. If sulphur 
dioxide is passed into a bottle containing nitrogen te- 
troxide, N 2 4 , the red fumes will soon disappear because 
the tetroxide has been deprived of a portion of its oxygen 
and converted into the dioxide, thus : — 

2 S0 2 + N 2 4 = 2 S0 3 + N 2 2 . 

Likewise a current of sulphur dioxide passed into a 
solution of potassium dichromate, or permanganate, will 
deprive them of a portion of their oxygen, changing the 
first to a compound, green in color, and rendering the 
second colorless. It will be important to remember this 
property on account of its relation to the manufacture of 
sulphuric acid, to be shown later. 

18. Uses of Sulphur Dioxide. — These have already been 
mentioned. It is used frequently as a disinfectant or 
fumigant, and for bleaching silk and straw goods. Evap- 
orated fruits, especially apples and peaches, owe their 
white, almost natural, color to the bleaching effects of 
sulphur dioxide, which is allowed to flow over the fruit 
as it is put into the evaporator. Its most extensive use is 
for making sulphuric acid. 

19. Sulphur Acids. — Sulphur forms several acids with 
hydrogen and oxygen, not all of which are important. 
The best known is sulphuric, H 2 S0 4 , also called oil of 
vitriol. 

Experiment 123.* — Arrange three flasks as shown in Fig. 47, one 
for the generation of sulphur dioxide by the treatment of copper with 

* If it is found necessary to use simpler apparatus, fill a flask with sul- 
phur dioxide, and introduce into it a swab of cloth upon the end of a glass 
j rod, moistened with nitric acid. Soon, brown fumes will begin to appear, 



180 



MODERN CHEMISTRY 



sulphuric acid, another containing nitric acid and copper turnings for 
the preparation of nitrogen dioxide, and a third containing water to 
be converted into steam. Connect with a large flask, D } which has a 
fourth tube to allow the entrance of air. 




Fig. 47. 

When the nitrogen dioxide enters the flask, D, containing air, it 
combines with the oxygen, forming the tetroxide, thus : — 

N 2 2 + 2 = N 2 4 . 

Immediately the sulphur dioxide attacks this compound of nitrogen, 
taking away two atoms, reducing it to the dioxide again, thus : — 

N 2 4 + 2 S0 2 = N 2 2 + 2 S0 3 . 

The dioxide thus formed, with the oxygen of the air, again combines 
to form the tetroxide, and so serves as a carrier of oxygen from the 
air to the sulphur dioxide. 

Next, the sulphur trioxide combines with the water introduced in 
the form of steam, producing sulphuric acid, thus : — 

S0 3 + H 2 = H 2 S0 4 . 



showing that the nitric acid is being decomposed and the sulphur dioxide 
converted into the trioxide. Now add a few cubic centimeters of water, 
and shake. The flask will contain a dilute solution of sulphuric acid. 



SULPHUR AND ITS COMPOUNDS 181 

20. To test the Acid prepared. — Put a little of the 
acid into a test-tube and add 1 or 2 cc. of a solution 
of barium chloride. If a white precipitate forms, which 
is not soluble in nitric or hydrochloric acid, or both to- 
gether, sulphuric acid is indicated. 

21. The Manufacture of Sulphuric Acid. — This acid is 
now prepared in immense quantities. The United States 
and Great Britain each produce annually about one million 
tons, and Germany is not far behind. Formerly, sulphur 
was used to prepare the sulphur dioxide for the manufac- 

i ture of this acid, but, as stated above, the attempt to con- 
trol the entire output of the Sicilian mines raised the price 
to such an extent that sulphuric acid manufacturers sought 
other sources, and finally discovered the present method. 
The pyrite is roasted in the presence of plenty of air, and 
the following reaction takes place : — 

2 FeS 2 + 11 O = Fe 2 3 + 4 S0 2 . 

22. These fumes are conducted into large chambers 
lined with sheet lead, into which jets of steam are con- 
stantly sprayed, together with nitric acid vapors, obtained 
by treating sodium nitrate with sulphuric acid. The 
reactions that take place in these lead chambers are the 
same as already described. The amount of nitric acid 
necessary is very small, and theoretically might be used 
indefinitely, but practically it is gradually carried by the 
draughts of air into the flues and must be replaced. 

The sulphuric acid thus prepared collects upon the floors 
of the rooms, — which are so large that a dancing party 
of a hundred couples could easily be held in them, — and 
is called chamber acid. It is only moderately strong, and 
is next evaporated in leaden vessels until a specific gravity 
of a little over 1.7 is reached, when it begins to attack the 



182 



MODERN CHEMISTRY 



lead. It is next transferred either to glass or platinum 
retorts, and concentrated until it reaches a density of 




Fig. 48. — Apparatus for condensing Sulphuric Acid. 

about 1.85. Fig. 48 shows the method of concentrating 
by glass retorts. The jars outside receive the nitric acid 
and other impurities contained in the sulphuric acid. 

23. Characteristics of Sulphuric Acid. — It is a colorless, 
sirupy liquid ; it received the name oil of vitriol on this 
account, and because it was made from green vitriol, ferrous 
sulphate. It is not a volatile acid, and, unlike nitric or 
hydrochloric, gives off no odor. It is very heavy and very 
corrosive. Organic matter exposed to it is charred black, 
as already noticed. It has great affinity for water, so 
much so that a beaker two-thirds filled with strong acid 
will in a few weeks, if left exposed to the air, absorb 
enough moisture to cause the beaker to overflow. Like- 
wise when strong acid is added to water, or vice versa, 
the mixture becomes very hot, reaching nearly 100° C, 
owing to the strong affinity of the two for each other. 

24. It is upon this principle that sugar, C 12 H 22 O n , is 
charred. The hydrogen and oxygen, being sufficient to 
form eleven molecules of water, are abstracted, and the 
carbon remains behind as a black mass. Upon the same 



SULPHUR AND ITS COMPOUNDS 183 

principle depends its use as a drying agent for various 
gases. They are made to bubble up through a bottle of 
strong sulphuric acid, and by this means lose their 
moisture. 

25. Uses for Sulphuric Acid. — It will be concluded from 
the vast quantities manufactured that sulphuric acid is a 
very important article of commerce. It is the most useful 
of acids, and almost all the others are dependent upon 
it for their preparation. In the manufacture of soda 
crystals, Na 2 C0 3 , by the Leblanc process (see page 211), 
sulphuric acid is indispensable. This salt, Na 2 C0 3 , is the 
basis for all soap manufacture as well as for glass, baking 
powders, etc. We can thus see the commercial importance 
of sulphuric acid. 

26. Another very extensive use is in the manufacture 
of artificial fertilizers from bones. When they have had 
the bone oil and gelatine removed, and as bone-black are 
no longer valuable for clarifying sugar, the bones are 
treated with sulphuric acid. This converts the phosphates 
present into a soluble form that may be used by plants. 
Sulphuric acid is also used in the manufacture of such 
explosives as nitroglycerine and gun-cotton, for making 
glucose, and in some of the processes of electroplating and 
electrotyping. 

27. Other Acids of Sulphur. — Sulphurous Acid, H 2 S0 3 . — 

IThis acid has already been mentioned, as well as its 
method of formation and its instability. 
We also have 

Hyposulphurous, H 2 S0 2 . 

Fuming Sulphuric, H 2 S 2 7 , 

which is really ordinary sulphuric acid, charged with 
sulphur trioxide, S0 3 . 



184 MODERN CHEMISTRY 

28. Thiosulphuric Acid, H 2 S 2 3 . — This last is of some 
interest because it is the basis of the salts known as 
thiosulphates, the best known of which is sodium thiosul- 
phate, Na 2 S 2 3 . From a mistaken idea of its composition 
sodium thiosulphate was first named hyposulphite, and is 
still commonly sold under that name. This is the photog- 
rapher's "hypo." 

SUMMARY OF CHAPTER 

Sulphur — Where found. 

Forms in which it occurs. 
Sources of commercial supply. 
Methods of purification. 
Characteristics of sulphur. 

Various forms — How prepared. 
How different. 
Uses of sulphur. 
Compounds. 

With hydrogen — Two names for the gas. 
Occurrence. 
Method of preparing. 

Characteristics of, and proof by experiment. 
Use of. 
With oxygen — Names and formulae. 
Preparation of the more important. 
Comparison of method with that of making hydrogen. 
Characteristics of S0 2 . 
Uses. 
With hydrogen and oxygen. 

Most important — Commercial name. 

How manufactured — Explanation of the chemical 

changes involved. 
Characteristics of H 2 S0 4 . 
Uses. 



CHAPTER XIV 
SILICON AND ITS COMPOUNDS —GLASS 

Silicon: Si = 28 

1. Abundance. — Silicon is never found free, but in 
the form of compounds is the most widely distributed 
and most abundant of the non-metallic elements except 
oxvgen and carbon. Sand, an oxide of silicon, Si0 2 , is 
familiar to all; quartz, crystallized or massive, including 
the agate, amethyst, opal, and other stones, is another 
variety of the same substance. All soils contain it to a 
greater or less extent, and it is taken up by plants and 
enters into their structure. Combined with sodium, cal- 
cium, magnesium, aluminum, and other metals it forms 
silicates which are very abundant. In this class may be 
placed granite, mica, and many other substances. 

2. Character of Silicon. — Silicon has been prepared in 
such limited quantities that not a great deal is known 
about it. It occurs in three forms, the amorphous and 
the crystallized or transparent being the two most impor- 
tant. At high temperatures it combines readily with 
oxygen or with carbon dioxide, forming the dioxide. 

Compounds of Silicon 

3. As already stated, silica or silicon dioxide, Si0 2 , 
is the most abundant compound. In the crystallized 
form it is often called rock crystal, and is found 
in hexagonal prisms, often more or less modified. Owing 

185 



186 MODERN CHEMISTRY 

to the presence of foreign substances, silica often assumes 
a variety of colors, and is known as rose quartz, smoky 
quartz, etc. It is very hard, being seven in the scale, is 
brilliant, highly refractive when cut, and is often used for 
ornaments in imitation of diamonds. It melts at about 
2000° C, and is soluble in alkalies as well as in hydro- 
fluoric acid. 

4. It is from the fact above mentioned that siliceous 
incrustations occur about many geysers. These springs 
are alkaline in character, and at the high temperature 
present beneath the surface dissolve considerable quanti- 
ties of silica ; when the water becomes cold and exposed 
to the action of the air, it is not able to hold the silica, 
and this is deposited upon any bodies on which the water 
may fall. The power of alkalies to dissolve silica may 
often be observed in the laboratory, where bottles contain- 
ing ammonia, caustic soda and potash, sodium carbonate 
solutions, etc., become etched or rough on the inside, and 
the glass stoppers so tight as to render their removal an 
impossibility. 

5. The Silicates. — Theoretically, silica is the anhy- 
dride of silicic acid ; that is, 

Si0 2 + 2H 2 = H 4 Si0 4 . 

But water added to the oxide in this case produces no 
reaction. The silicates, however, are based upon this acid. 
They are abundant, and many of them are very complicated 
in composition. As silicic acid is tetrabasic, the hydrogen 
may be replaced by a variety of elements, even in the 
same molecule ; thus, we might have 

NaAlSi0 4 : Sodium Aluminum Silicate. 
CaMgSi0 4 : Calcium Magnesium Silicate. 



SILICON AND ITS COMPOUNDS — GLASS 187 

NaKCaSi0 4 : Sodium, Potassium Calcium Silicate. 
H 8 Mg 6 Fe 7 Al 2 Si 8 18 : Mica, etc. 

6. Preparation of Silicic Acid. — Silicic acid may be pre- 
pared from " water glass," that is, silica dissolved in boiling 
caustic soda, or potash, by adding a little strong hydro- 
chloric acid till the solution is no longer alkaline. Then 
a jellylike mass will be precipitated, which is silicic acid. 
By filtering this out and igniting when dry we again 
obtain the oxide. 

Experiment 124. — Let the student thus prepare some soluble 
u water glass " and the silicic acid from it. 

7. Though silica is insoluble in water and has such a 
high melting point that only such temperatures as that 
secured by the oxyhyclrogen blowpipe or the electric fur- 
nace will fuse it, still, if mixed with sodium carbonate and 
strongly heated in a blast lamp for a few minutes, it is 
converted into a soluble form, sodium silicate, and may 
then be readily taken up by water. 

8. Glass. — This is an artificial silicate that has been 
manufactured in some form or other for probably 4000 
years. Several of the nations of antiquity were famous 
for their wonderful glasswork ; in beauty of coloring, 
their achievements have probably not been surpassed in 
modern times. But the applications of glass are now so 
varied and so adapted to the necessities of life, as well as 
to the luxuries, that it would seem impossible to do with- 
out it. Every year sees the manufacture of hundreds of 
millions of bottles, and tons of other kinds of glassware ; 
and the art of glass blowing and working has reached such 
a high state of perfection that glass objects, from their 
nature almost inconceivable, are now of frequent manu- 
facture. 



188 MODERN CHEMISTRY 

We have seen above that if silica is fused with sodium 
carbonate, a new compound is formed which is quite 
soluble. If, however, we mix calcium carbonate, or chalk, 
with the silica, together with the sodium carbonate, and 
fuse the mixture, we then obtain a double silicate of 
sodium and calcium that is quite insoluble in water and 
in all acids, except hydrofluoric. 

9. Varieties of Glass. — There are many varieties of glass. 
As potassium salts are so closely related to those of sodium, 
it is obvious that potash could be used instead of soda. 
In fact, glass was first made entirely in this way. Nearly 
all the best chemical and physical apparatus is still made 
from potash salts, and this variety is known as Bohemian 
glass. It is much harder to melt than glass made from 
sodium carbonate. 

10. If an oxide of lead is used with the silica and potash, 
we obtain a glass that is very soft and easily worked, 
known as flint glass. It has a very high refractive power, ; 
and on this account is used for telescopes and all kinds of 
optical instruments. In the purest form it is known as 
strass or paste, and from this are made the so-called paste 
diamonds. These are so lustrous and highly refractive 
that, except in hardness, it is difficult for any but experts 
to distinguish them from the genuine article. 

11. Ordinary glass, known as crown glass, from which 
windows and the great majority of glass utensils are made, 
is a silicate of lime and sodium, as already described. 
Ordinary sand contains a considerable amount of iron 
in the form of an oxide. This gives to the glass used 
for ordinary bottles and for all the cheaper grades the 
well-known greenish color, which, however, may be re- 
moved by the addition of a small quantity of manganese 
dioxide. 



SILICON AND ITS COMPOUNDS — GLASS 189 

12. Much of the plain glassware used at present is molded 
just as any casting would be in an iron foundry. Window 
glass is first blown into a long cylinder ; this is cut open 
and flattened while still hot by means of heavy rollers. 
Plate glass for large windows and heavy mirrors is cast. 
The molten glass is poured upon a table of the desired 
size, allowed to cool, and the surface afterwards ground 
.and polished. 

13. All glass articles must be carefully annealed, other- 
wise they would be so brittle as to have little value. The 
glass, as soon as shaped, is placed in an oven, and during 
several days is cooled so slowly that the molecules have 
time to adjust themselves to stable positions. Indeed, so 
well is this annealing done that glass vessels are made 
for use in chemical work that may be heated strongly and 
plunged into cold water immediately without danger of 
breaking. 

SUMMARY OF CHAPTER 

Silicon. 

Abundance of it in nature. 

Some familiar forms. 
Compounds of silicon. 

The oxide — Some common forms. 
Characteristics of. 
Glass — Importance of. 

What glass is. 

Kinds of glass. 

How different in properties. 

Making of window glass and other forms. 

Annealing of glass articles. 



CHAPTER XV 



PHOSPHORUS AND ITS COMPOUNDS 



Phosphorus : P = 31 



1. Occurrence. — Phosphorus has been known for about 
two and a quarter centuries, but it is only since 1833 that 
it has had any real practical value. Owing to its strong 
affinity for oxygen it is never found free, but in the form 
of compounds it is very widely distributed. It is a con- 
stituent of many rocks, and, from their decomposition, also 
of soils. From this source plants take it up and store it 
away in the seeds and fruits ; plants, being used as foods, 

transfer it to animals, where it 
is found in the nerve centers 
and bones. 

2. Manufacture of Phospho- 
rus. — It is obtained almost 
altogether from bones. These 
are put into retorts and heated, 
much as coal is for the prepa- 
ration of illuminating gas. The 
volatile products are thus driven | 
off and their valuable portions 
condensed. The bones are re- 
duced to what is known as 
bone-black, or, if not desired for 
clarifying sugar, to bone-ash. 
To this sulphuric acid is added, which converts the cal- 
cium phosphate in the bones into a salt that is soluble in 

190 




Fig. 49. - 



- Manufacture of Phos- 
phorus. 






PHOSPHORUS AND ITS COMPOUNDS 191 

water. This is dissolved out and the solution evaporated 
to dryness, then mixed with carbon and strongly heated. 
The phosphorus is thus set free ; it distills out and is 
condensed under water and molded into sticks. (See 
Fig. 49.) i2, i2, are the retorts into which the mixture 
of charcoal and phosphorus compounds are put ; F is the 
furnace, and W, TF, the water tanks where the phosphorus 
is condensed. The process is very deleterious to health, 
the fumes from the retorts often producing dangerous 
ulcerations of the jawbones, a disease which is practically 
incurable. 

3. Characteristics of Phosphorus. — Phosphorus is a 
very pale, amber-colored, translucent solid, somewhat waxy 
in appearance. When exposed to the air it almost imme- 
diately begins to give off luminous fumes having a faint 
garlic odor, and in the course of a short time takes fire. 
A little friction will readily ignite it, hence it should be cut 
under water. Burns from it are very serious and require 
weeks to heal. If heated to 210° out of contact with the 
air, it changes to an ailotropic form, known as red or 
amorphous phosphorus. Unlike the ordinary phosphorus, 
this is not poisonous, does not readily take fire, is not solu- 
ble in carbon disulphide, and does not glow in the dark. 

Experiment 125. — To show the ready combustibility of phos- 
phorus when finely divided. Dissolve a small piece of phosphorus, 
half as large as a pea, in a little carbon disulphide. Pour the solution 
upon a piece of filter or blotting paper, and let it dry. Notice how 
quickly it ignites. 

Experiment 126. — To show that phosphorus will burn under 
; water. Put into a small bottle about 1 g. of potassinm chlorate, add 
a few small pieces of phosphorus, and cover with water. By means of 
a pipette or funnel tube introduce beneath the water into contact with 
the potassium chlorate a little siilphnric acid. Notice that the phos- 
rphorus begins to burn. Explain. 



I 



7 



192 MODERN CHEMISTRY 



Experiment 127. — To show the affinity of phosphorus for chlo- 
rine, bromine, and iodine. Put a small piece of phosphorus into a 
deflagrating spoon and introduce it into a jar of chlorine. What 
happens in a few moments? Cut a thin slice of phosphorus, and 
upon it place a crystal of iodine. Notice that the phosphorus is soon 
ignited. Try a drop of bromine in the same way. 

4. Uses for Phosphorus. — About 3000 tons of phos- 
phorus are manufactured annually, most of which is used 
in preparing matches. Small quantities also are employed 
ill making poisons. Matches were first made in Austria 
by tipping small pine sticks with sulphur to which a little 
phosphorus had been added. This method was employed 
for a good many years, but the sulphur has now been 
largely replaced by other substances rich in oxygen, such as 
potassium chlorate, saltpeter, etc., together with paraffine. 

5. Matches are now made entirely by machinery, and 
with wonderful rapidity. The wood, being sawed into 
convenient lengths, is pressed against knives, which split 
it up into the proper size for matches. These are dipped 
into paraffine, then tipped with a paste made of a little 
glue containing phosphorus and the other ingredients 
already mentioned, together with some coloring matter. 
After drying they are packed in boxes. In this way 
a single machine will make and pack several million 
matches in a day. In the case of safety matches the 
phosphorus is placed in the prepared surface upon the, 
box, and the matches can be ignited only by friction on 
this surface. 

6. Compounds of Phosphorus. — One of the most inter- 1 
esting of these is hydrogen phosphide, PH 3 . It is also ; 
called phosphine and phosphoreted hydrogen. It is readily 
evolved when phosphorus is heated in a solution of any 
strong alkali, such as caustic soda or potash. 



PHOSPHORUS AND ITS COMPOUNDS 



193 




Experiment 128. — Suitable for class-room. Into a small flask put 
about 50 cc. of strong caustic soda or potash solution, and add several 
small pieces of phos- 
phorus. Pour in about 
a cubic centimeter of 
ether, and close the 
flask quickly with a 
cork and long delivery 
tube. Support the 
flask upon a ring-stand, 
as shown in the figure, 
and heat moderately. 
Presently smoky-look- 
ing fumes will fill the 
flask, and then the 
bubbles issuing from 
the mouth of the de- 
livery tube will take 
fire spontaneously. If 

the room is free from draughts of air, beautiful rings of smoke, grow- 
ing gradually larger, will float upward. Notice the vortex motion of 
:he rings. The ether is introduced to expel the air before any phos- 
phine is generated ; the heat should be regulated so as not to allow 
:oo rapid an evolution of gas, otherwise the rings will follow in such 
rapid succession as to break one another. What is the odor of the 
tas? Color? 

7. Oxides of Phosphorus. — Pentoxide, P 2 5 , and Tri- 
Dxide, P 9 Oq. The first of these has been seen on various 
occasions : when phosphorus was burned in oxygen, in 
preparing nitrogen, etc. The dense white fumes noticed 
consisted mainly of phosphorus pentoxide. This com- 
pound is always obtained when phosphorus is burned in a 
plentiful supply of oxygen. When the amount is limited, 
br when the combustion is slow, phosphorus trioxide is 
obtained. The pentoxide is a white solid which has great 
imnity for moisture, and if dropped into water combines 
with it with a hissing sound as of a hot iron in cold water. 



194 MODERN CHEMISTRY 

8. Acids of Phosphorus. — The two oxides named above, 
like the corresponding oxides of nitrogen, are the anhy- 
drides of certain acids, thus : — 

P 2 3 + 3 H 2 = 2 H 3 P0 3 . . . Phosphorous Acid 
P 2 0- 6 + 3 H 2 = 2 H 8 P0 4 . . . Phosphoric " 

The latter is the more important. It will be noticed that 
its molecule contains three atoms of hydrogen, all of which 
may be replaced by a metal. Such acids are called tribasic. 
Phosphoric acid is a white crystalline substance, which 
may be prepared by treating bone-ash with sulphuric acid. 
At high temperatures it will give up a part of the water 
that was taken in its formation and yield metaphosphoric 
acid, HP0 3 , which is monobasic. The reaction may be 
shown thus : — 

H 3 P0 4 + heat = HP0 3 + H 2 0. 

This is frequently sold under the name glacial phosphoric 
acid. 

9. Compounds with Phosphoric Acid. Phosphates. — The 

most common of these is calcium phosphate, Ca 3 (P0 4 ) 2 , 
found in the bones. Immense deposits of this are found 
in Florida, where it is mined and used as a fertilizer in 
various parts of the world. From the fact that all grain 
plants absorb the soluble phosphates from the soils, unless 
these salts are replaced in some way the land rapidly loses 
its productive power. A considerable portion of the phos- 
phates in the grain fed to animals is thrown off in the ex- 
crement and is returned to the soil in this way. 

10. Immense quantities of bones are reduced to animal 
charcoal, and then, by treatment with sulphuric acid, con- 
verted into soluble phosphates and returned to the soil 



PHOSPHORUS AND ITS COMPOUNDS 195 

in this way as artificial fertilizers. Another source of 
considerable supply is from the reduction of phosphorus- 
bearing iron ores by the Thomas-Gilchrist process ; and a 
matter of interest in this connection is that the calcium 
phosphate thus obtained is already in the soluble form and 
needs no further treatment. 



SUMMARY OF CHAPTER 

Phosphorus — Occurrence. 
Source of supply. 
Method of preparing phosphorus. 

By-products and their uses. 
Characteristics. 

Two forms of phosphorus. 

Compare with forms of sulphur. 
Experiments to illustrate characteristics. 
Uses. 

Method of making matches. 
Chemicals used. 
Compounds. 

With hydrogen. 

How prepared. 
With oxygen. 

Names and formulae. 
How prepared. 
With hydrogen and oxygen. 
How related to the oxides, 
Salts formed from these acids. 
Uses. 



CHAPTER XVI 

AVOGADRO'S LAW — ATOMIC WEIGHTS — PROBLEMS 

1. Avogadro's Law. — This law, or hypothesis, was for- 
mulated by the Italian physicist and chemist, Avogadro, 
and afterward, independently, by the Frenchman, Ampere. 
It may be stated thus : — 

Equal volumes of all substances in the gaseous condition 
under the same pressure and temperature contain the same 
number of molecules. To illustrate, suppose a liter of 
hydrochloric acid gas contains a billion molecules, then 
a liter of nitrous oxide, or of any other gas, would also 
contain a billion molecules. 

2. Proof of this Theory. — No absolute proof of this- 
law has ever been given, but many facts seem to favor 
such a theory. For example, we have seen that all gases 
expand and contract in the same ratio under the influence 
of heat and pressure. As expansion and contraction mean 
simply a change in the distance which separates the mole- 
cules from each other, this being greater when the body 
is heated, and less when cooled, it would seem that bodies - 
could expand alike only if composed of the same number 
of molecules, or if containing what means the same thing, 
the same number of intermolecular spaces. 

3. Ratio of Molecular Weight to Specific Gravity. — It 
has been found also that there is a constant ratio existing 
between the molecular weight of a gaseous body and its 
specific gravity ; that is, if we divide the molecular weight 
of any gas by its specific gravity, we always obtain prac 

196 • 



ATOMIC WEIGHTS 197 

tically the same quotient. This ratio is about 28.88. 
Thus, the molecular weight of carbon dioxide is 44, its 
specific gravity is 1.524, the ratio of 44 to 1.524 is 28.87 ; 
carbon monoxide has a molecular weight of 28, specific 
gravity of 0.967, the ratio is 28.94. 

Exercise. — To apply this fact, suppose a given volume of nitrous 
oxide, X 2 0, weighs 1.52 grams, and the same volume of hydrochloric 
acid gas weighs 1.27 grams. It is evident that the weight of any 
volume of gas divided by the weight of one molecule would give the 
number of molecules in that volume. Thus : — 

wt. of 1 1. N 9 i a- /-. • i i*I 

2— — ll0 . niol. >s 9 m 1 liter. 

wt. of 1 mol. 

and wt. of 1 1. HC1 = moL HC1 - n x j.^ 

wt. of 1 mol. HC1 

We can readily find the weight of a liter of each of these gases, 
and also the molecular weight of each, but the first is in grams and the 
second in microcriths^ that is, so many times as heavy as a hydrogen 
atom ; but unfortunately we have no means of knowing how many 
microcriths in a gram, hence we cannot perform the division indicated 
above nor assign to the quotient any concrete name. If we make 
the division, however, we find that the quotient is always practically 
the same ; that is, 

wt. of i vol. y,o = 28 88 = no moL N in 1 voL 

wt. of 1 mol. X 2 

wt. of 1 vol. HC1 = 28b88 = no< moL HC1 . j yol 
wt. of 1 mol. HC1 

Hence, as the ratio in each case is the same, in accordance with the 
axioms of geometry, 

no. mol. X 2 in 1 vol. = no. mol. HC1 in 1 vol. 

Putting this law into the form of a proportion, it would read : — 

mw m'w' , , , 

■ — = , or mw : mw : : s : s , 

s s' 

in which mw and m'w' represent the molecular weights of any two 
gases, and s and s' their specific gravities. 



I 



198 MODERN CHEMISTRY 

4. Application of this Law. — The truth of Avogadro's 
Law having been accepted long ago, it is now made use of 
in determining the molecular weights of new compounds. 
Having found by actual work the weight of 1 liter of the 
gas, and knowing the weight of 1 liter of air, the specific 
gravity is found. Then, substituting in the formula, 

— = 28.88, 
s 

we can easily find the value of mw. 

5. Finding Atomic Weights. — This law is further used 
in determining the atomic weight of a newly discovered 
element. 

Let m represent this element, and suppose we are attempting to 
find the atomic weight by studying some compound of it with oxygen. 
We should find the weight of a molecule of the oxide as shown above. 
Suppose this is found to be 28. Next, by chemical analysis we should 
determine what per cent of the compound is the new element, m. 
Suppose the analysis shows this to be 42.86 per cent, then we should 
have this proportion : — 

mol. wt. : wt. of m in the mol. : : 100 per cent : per cent of m; 
or, 28 : x : : 100 : 42.86. 

100 x = 42.86 x 28. 
x = 12. 

In the same way we should determine the weight of the element, m, 
in a molecule of a number of other compounds containing it; then, 
the one having the least amount would be taken as a compound con- 
taining but one atom of the element, and the value of x in that com- 
pound would represent the atomic weight. To illustrate, suppose in 
this way we find in our analyses, and subsequent determinations, that 
m or x is equal to 24, 12, 36, 120. The second, 12, being the smallest 
amount found in any compound, would be accepted as the atomic 
weight. This would not be absolute proof, however, as later another 
compound might be discovered which contained a smaller amount of 
m, in which case that smaller amount would be taken as the atomic 
weight.. 



ATOMIC WEIGHTS 



199 



6. Constitution of the Molecules of Elements. — How 
many atoms are there in a molecule of an elementary sub- 
stance, like oxygen, hydrogen, etc. ? In writing some of 
the reactions in the earlier part of this book the molecules 
were shown as having two atoms. With some exceptions, 
this is true ; that is, a molecule of hydrogen, oxygen, 
chlorine, and of many other elements contains two atoms. 
How do we know this? A proof in the case of one or 
two elements will illustrate for the others. 

We have seen that when hydrogen and chlorine are caused to unite, 
they form hydrochloric acid. It is found also by further experiment 
that in uniting thus the volume is not decreased ; that is, if we put 
together a liter of chlorine and one of hydrogen, after causing them 
to combine, we have 2 liters of hydrochloric acid. Perhaps it will be 
clearer, stated in the form of an equation, thus : — 

1000 cc. of CI -f 1000 cc. of H = 2000 cc. of HG1. 

Now, according to Avogadro's Law, there would be the same number 
of molecules in a liter of chlorine as of hydrogen or of hydrochloric 
acid. Dividing the entire equation through by this common factor, 
the number of molecules of chlorine in 1 liter, or 1000 cc, we should 
have 

1 mol. of CI. + 1 mol. of H = 2 mol. of HC1. 





Two Molecules. 



Chemical analysis shows that in hydrochloric acid the hydrogen and 
chlorine are united in the ratio of 1 to 35.5, or one atom of each, 
as represented by the formula HC1, or by the figure. 



200 MODERN CHEMISTRY 

It is evident, therefore, that two molecules of hydrochloric acid 
contain two atoms of hydrogen and two of chlorine, and as we only 
had one molecule of each of these elementary gases, each of those 
molecules must have contained two atoms. In a similar way we 
would prove for bromine, fluorine, oxygen, and others. 

7. Most Molecules Diatomic. — Such molecules as these 
are called diatomic. There are a few, sodium, potassium, 
cadmium, mercury, and zinc, whose molecules contain 
only one atom, and such are called monatomic. Their 
molecule is, therefore, identical with the atom. Only one 
triatomic elementary molecule is known, and that is the 
allotropic form, ozone. A few, like phosphorus and arsenic, 
are tetr atomic ; that is, the molecule is made up of four 
atoms. 

8. Application of this Fact. — It often becomes neces- 
sary in chemical problems to know the weight of a liter 
of a gas. This may very easily be found, but we must 
first know its vapor density ; that is, its density compared 
to hydrogen. With the elementary substances this is, 
as a rule, the same as the atomic weight ; for example, 
the atomic weight of hydrogen is 1, the molecular weight 
is 2 ; the atomic weight of nitrogen is 14 ; the molecular 
weight 28. Hence, whether we take the atomic weight of | 
nitrogen, or its molecular weight and divide by the molec- 
ular weight of hydrogen, we obtain the same results. 
Then, as the hydrogen molecule weighs two, we find the 
vapor density of any other substance by dividing its 
molecular weight by 2. Thus : — 

1 mol. N 2 weighs 2 x 14 + 16 = 44 

1 " H " 2x1 = 2 

1 " N 2 " 44 -T- 2 times as much as 1 mol. H 



ATOMIC WEIGHTS 201 

Again, — 

1 mol. CO weighs 12 + 16 = 28 

1"H " 2x1=2 

1 " CO " 28 -=- 2 times as much as 1 mol. H 

Therefore, vapor density of CO is 28 -s- 2 = 14. 

Thus find the vapor density of C0 2 , N 2 3 , O, HC1, S0 2 , 
CI, N. 

9. To find Weight of One Liter of Any Gas. — Having 
found the weight of a gas compared to hydrogen (its 
vapor density), it is only necessary to multiply the weight 
of 1 liter of hydrogen by this figure. A liter of hydro- 
gen has been found to weigh .0896 g., a number which 
should be remembered. Suppose now we desire to find the 
weight of a liter of carbon monoxide, CO. Above Ave 
found its vapor density to be 14. Then, as a liter of 
hydrogen weighs .0896 g., one of carbon monoxide will 
weigh 14 x .0896, or 1.2544. 

Thus find the weight of 1 liter of the gases whose densi- 
ties were found above. Also of N 2 0, NH 3 , H 2 S. 

10. The Formulae of Compound Bodies. — We have 
learned that the formula of a compound is a short method 
of expressing its composition. It may be of interest to 
know how to determine the formula of a compound. The 
substance is first carefully analyzed, and the percentage 
composition determined. 

Suppose we have in mind a compound which analysis shows con- 
sists of carbon and oxygen, 27.27 per cent of the former, and 72.73 per 
cent of the latter. We should next w T eigh a liter of it ; suppose we 
I find this to be 1.9712 g. As a liter of hydrogen weighs .0896 g., the 
unknown gas is 1.9712 -^ .0896, or 22 times as heavy. 

We have seen that the molecular weight is twice the vapor density, 
J then the weight of the molecule would be 2 x 22, or 44. Now, as the 



202 MODERN CHEMISTRY 

carbon is 27.27 per cent of this, it equals .2727 of 44 = 11.9988, and the 
oxygen, 72.73 per cent, or its weight in the molecule is 72.73 per cent of 
44 = 32 +. Previous experiments have shown that the atomic weight 
of carbon is 12, hence the w r eight found above, 11.9988, practically 
corresponds to one atom, and that would be the amount of carbon in 
the compound. In the same way as the atomic weight of oxygen is 
known to be 16, the amount found in the compound, 32, would indi- 
cate two atoms. The substance in question, therefore, . would contain 
carbon, 1 atom, oxygen, 2 atoms, and would be carbon dioxide, for- 
mula C0 2 . 

Problems. — 1. A liter of a certain gas weighs 0.8064. It consists 
of hydrogen ^ and oxygen |. Find its vapor density, molecular weight, 
and the formula. 

2. A gas consisting of carbon and oxygen has 42.86 — per cent of 
the former, and 57.14 -f per cent of the latter. If 1 liter of it weighs 
1.2544, what is its formula? 

3. What per cent of turpentine, C 10 H 16 , is carbon? Hydrogen? 

4. The vapor density of a body is found to be 50.5. If analysis 
shows that 2.359 g. of it contain 1.12 g. of oxygen, how many atoms 
of oxygen are there in the formula representing the substance ? 

5. What is the molecular w 7 eight of a certain substance if 50 g. 
of it contain 32.65 g. of oxygen, knowing that there are four atoms of 
oxygen in the molecule of the substance? 

6. Find the percentage composition of nitric acid. 

SUMMARY OF CHAPTER 

Avogadro's Law — Statement of the law. 
Illustration. 
Proof of the law. 

a. As seen in effects of heat. 

6. Ratio of molecular weight to specific gravity. 
Value of the law. 

a. Finding atomic weights — Illustration. 

b. Constitution of molecules — Proof. 
Meaning of terms monatomic, etc. 

Problems. 

Method of finding weight of a liter of any gas. 
How to determine the formula of a compound. 



CHAPTER XVII 

THE METALS— PERIODIC LAW 

1. Metals and Non-metals. — It has been customary to 
divide the elements into two great classes, the metals and 
non-metals, of which the former includes by far the greater 
number. This classification, however, is based largely 
upon the external characteristics or appearance rather 
than upon the chemical deportment. In appearance the 
metals have a peculiar luster, known as the metallic luster, 
considerable density, with few exceptions have high melt- 
ing points, and are electro-positive in character. As a 
rule, their oxides are not anhydrides, and yet there are 
many exceptions to this statement, for we find various 
compounds of tin, arsenic, antimony, chromium, aluminum, 
etc., in which these metals seem to serve as the acid-form- 
ing element. And some even possess more chemical char- 



Non-metals 


Metals 


Their oxides form acids, as for 


Their oxides form bases as : 


example : 




N 2 3 . . . HN0 2 , 


CaO . . . Ca(OH) 2 , 


S0 2 . . . H 2 S0 3 , 


Na 2 . . . NaOIL 


P 2 O s . . . HP0 3 ,etc. 


K 2 . . . KOH, etc. 


Many are gaseous. 


Most are solids. 


Many are transparent. 


All are opaque. 


Poor conductors of heat and 


Good conductors of heat and 


electricity. 


electricity. 



203 



204 



MODERN CHEMISTRY 



acteristics in common with the non-metals than with the 
metals. It must be concluded, therefore, that there is no 
clearly dividing line between the two classes. Neverthe- 
less, some distinctions in addition to those mentioned 
above may be noted. 

2. Tabular Classification. — It will be seen that the 
above division is almost purely an arbitrary one. At the 
present time it is customary to classify the elements into a 
number of groups in accord with what is known as the 
periodic law. 

Table of Elements 





I 


II 


III 


IV 


V 


VI 


VII 


VIII 


Period I 


H = l 
Li = 7 


Gl = 9 


B = ll 


C = 12 


N = 14 


= 16 


F = 19 




II 


Na 


Mg 


Al 


Si 


P 


S 


CI 




- ™{i 


K 

Cu 


Ca 

Zn 


Sc 
Ga 


Ti 

Ge 


V 

As 


Cr 

Se 


Mn 
Br 


Fe, Co, 
Ni 


•• Hi 


Rb 

Ag 


Sr 
Cd 


Y 

In 


Zr 

Sn 


Nb 
Sb 


Mo 
Te 


I 


Ru, Rh, 
Pd 


•• Hi 


Cs 


Ba 


La 


Ce 










•• Hi 


Au 


Hg 


Yb 
Tl 


Pb 


Bi 


W 




Os, Ir, 
Pt 


- ™{i 






Th 






U 







3. Recurrent Characteristics in the Table. — If the above 
table is studied in connection with the atomic weights of 
the elements, it will be seen that, reading from left to right, 



THE METALS — PERIODIC LAW 205 

they are arranged with reference to their weights. Thus, 
in the first period, we have 

Li = 7, Gl = 9, B = 11, C = 12, N = 14, O = 16, F = 19 ; 

in the second, 

Na=23, Mg=24, Al=27, Si = 28, P = 31, S=32,C1=35.5. 

4. In thus arranging them it was noticed that, starting 
with lithium, not until we reach the eighth element beyond, 
do we come to another, sodium, similar to lithium in char- 
acteristics ; and from sodium there are seven more before 
another is reached similar to this. From these observa- 
tions the above table was arranged, and though it is far 
from complete, wonderful results have come from it. We 
notice that in group VII, we have fluorine, chlorine, 
bromine, and iodine, four elements that we have found 
to have very similar properties. We shall hereafter find 
the same to be true of lithium, sodium, and potassium 
of the first group ; magnesium, calcium, strontium, and 
barium of the second, and so on. If we take these vertical 
columns or groups and compare their atomic weights, we 
notice some interesting facts. 

t • ry 1 

■ ~ „„l The atomic weight of sodium is exactly 
I halfway between the other two. 



K = 39 j 

Ca= 40 
Sr = 87 
Ba = 136 



The weight of strontium is practically 
the mean of the other two. 



P = 31 

As = 75 j- The same is true of the middle element. 

Sb = 120 



r 



206 MODERN CHEMISTRY 



The same is true of the middle element. 



s = 


32] 


Se = 


79 


Te = 


125 J 



5. If we study the compounds that the elements form, 
we shall find that those falling in the same group are 
strikingly similar in their chemical behavior. Thus, 
lithium, sodium, potassium, rubidium, and caesium in the 
first group are all univalent and form oxides with the 
general formula, M 2 0, in which M represents any metal 
of the group. Furthermore, they form no compounds 
with hydrogen. If we take the second group, they are 
all bivalent, forming oxides with the general formula, 
MO, as MgO, CaO, etc. They form no hydrogen com- 
pounds. The members of the third group are trivalent, 
as seen in their oxides, A1 2 3 , general formula M 2 3 . 
And so we might go on through the table. 

6. Vacancies in the Table. — It will be noticed that there 
are many vacant places, but it is an interesting fact that 
when the table was first worked out there were many 
others that have since been filled. And strange to say, 
from this table the author of the plan not only predicted 
that these very elements would be found, but even gave 
in a general way their characteristics, and in accordance 
therewith suggested names for them. In the same way, 
it is possible that many of the places now vacant will 
sometime be filled by elements as yet undiscovered. 

Note. — Some teachers may prefer to defer a close study of the Periodic Law until after 
the completion of the work with metals. 

SUMMARY OF CHAPTER 

Classes of the elements. 
. Characteristics of the two classes. 
Wherein different. 
Wherein alike. 



•> 



THE ALKALI METALS 207 

The Periodic Law. 

Recurrence of certain characteristics. 
Relation of atomic weights. 
Similarity of chemical behavior. 
Value of the law and table. 



CHAPTER XVIII 

THE ALKALI METALS 

Sodium : Na = 23 

1. Its Discovery. — Up to the year 1807 caustic soda 
and caustic potash had been regarded as elementary sub- 
stances ; by electrolysis, however, Sir Humphry Davy in 

1807 proved both of these substances to be compounds, 
and hydroxides of the metals sodium and potassium. 

2. Where found. — Sodium is very widely distributed, 
traces of it being found everywhere. On account of the 
strong affinity existing between it and water, it never 
occurs in the metallic state. Its most abundant com- 
pound is common salt, NaCl, which constitutes a large 
per cent of the solid matter found in sea water, salt 
lakes, and springs ; vast deposits of it, more or less pure, 
occur in many parts of the West as well as in other 
portions of the world. Sodium nitrate, NaN0 3 , is found 
in immense quantities in Chile and elsewhere. Other com- 
pounds occur in smaller proportions, but in some form or 
other sodium can be detected in every particle of dust that 
may be seen floating in the sunbeams. 

3. Reduction of Sodium from its Compounds. — Since 
the isolation of the metal by Davy, various other plans 

;^have been tried, but they are all modifications of the 



208 MODERN CHEMISTRY 

original. What is known as the Castner process is the 
one generally used at present. See Figure 51. 

A in the figure is a large iron 
«X^k^\_ vessel, B another, similar but 

^-— # A %es= \ smaller, inverted over A and 

fBl|: y^Sjj ) dipping into the fused caustic 

|8L____ il 1 jjf ( soda in the lower vessel. It is 

^^^^^^|^»^fe held in position by' insulated 

^fe^^|f^, j ' (7^\ supports not shown. Through 
M^B^^^^^l^J^) the bottom at D is inserted the 
\JL_^9^ """"" negative electrode, and B serves 
as the positive. When the cur- 
rent from the dynamo is passed through, the caustic soda 
is electrolyzed, B gradually fills with hydrogen which 
bubbles out underneath, while the metallic sodium col- 
lects upon the surface of the fused mass. In this way 
it is prepared for about two dollars a pound. 

4. Characteristics of Sodium. — It is a silvery white 
metal, so soft at ordinary temperatures that it may be 
molded with the fingers, about like stiff putty. At 
— 20° C, however, it becomes hard. It tarnishes so rapidly 
in the air that only for an instant after being cut can its 
true color be seen. It takes up moisture and carbon di- 
oxide from the air, forming first caustic soda, and after- 
ward sodium carbonate. In course of time a, piece of 
sodium left more or less exposed is entirely converted into 
amorphous sodium carbonate. It is usually preserved in 
naphtha or some similar light oil containing no oxygen. 

5. Sodium is soluble in liquid ammonia and forms with 
it a blue solution. Its properties are strongly alkaline. 
If heated and plunged into a jar of chlorine it burns vigor- 
ously, forming common salt. Thrown upon water it is 
immediately melted, owing to the heat generated by the 



THE ALKALI METALS 209 

strong chemical action, and the water is decomposed, as 
already shown in our study of hydrogen. If a burning 
match is touched to the sodium as it spins about on the 
water, the hydrogen will burn with a yellow flame, due to 
the vaporization of a small portion of the sodium. Upon 
moderately warm water the gas will take fire spontane- 
ously. If a small piece of sodium is dropped upon a 
moistened blotting paper, it is quickly ignited. If, when 
it begins to burn, the molten sodium is allowed to roll off 
and drop upon the floor, it will burst into many particles 
which will spin about, burning with the characteristic yel- 
low flame. 

Experiment 129. — Moisten a piece of blotting paper with water, 
to which a little phenolphthalein has been added. Drop a small 
piece of sodium upon the blotter. Notice the red titlntj^ft leaves as it 
slowly moves about from place to place. You have seen similar results 
in previous work. Let the molten globule of sodium roll off upon the 
floor and notice what happens. 

Compounds of Sodium 

6. Caustic Soda, Sodium Hydroxide, NaOH. — This com- 
pound is prepared by treating sodium carbonate, Na 2 C0 3 , 
in solution with lime-water. The reaction is 

Na 2 C0 3 + Ca(OH) 2 = 2 NaOH + CaC0 3 . 

The calcium carbonate, thus formed, is insoluble in water, 
hence is precipitated. The sodium hydroxide is drawn off, 
evaporated to dryness, purified, then fused and molded 
into sticks ; in this form it is put upon the market. It is a 
white solid, deliquescent, with strongly alkaline properties. 

7. Sodium Chloride, NaCl. — As already stated, this 
compound occurs very abundantly. In some places it is 
mined much as rock or metallic ores are mined. In other 



210 MODERN CHEMISTRY 

places, where the deposits are upon the surface, mingled 
with considerable quantities of sand and earthy matters, it 
is dissolved out and the strong solution evaporated. In 
some of our states wells are sunk into the deposits, and 
water pumped in to dissolve the salt. This is again drawn 
out and evaporated. In some places along the Mediter- 
ranean the sea water is pumped up and allowed to trickle 
down over brush or lattice work, whereby it is much con- 
centrated in strength, then this solution is evaporated to dry- 
ness in large shallow pans. It crystallizes in cubes, as may 
be seen if a strong solution is allowed to evaporate slowly. 

Sodium chloride, if chemically pure, is not deliquescent, 
but owing to impurities present that which is generally 
put upon the market soon becomes damp when exposed to 
the air. It is used very extensively in the manufacture of 
other important compounds of sodium ; also largely in our 
food. A part of it is said to be decomposed by the diges- 
tive fluids of the stomach and to form hydrochloric acid. 

8. Sodium Carbonate, Na 2 C0 3 . — This is a very impor- 
tant compound used in the manufacture of soap, glass, and 
for a variety of other purposes. In the early part of this 
century soda crystals, as this compound is often known in 
commerce, sold for over $300 a ton, while now the same 
quantity is worth scarcely $20. There are two general 
processes of manufacture. The simplest and the one most 
in favor at the present time is the — 

Solvay Process. — This consists of passing a current of ammonia 
into a strong solution of sodium chloride until it is saturated ; carbon 
dioxide is next forced in and with the ammonia forms ammonium 
bicarbonate. This reacts with the common salt, forming sodium 
bicarbonate. These processes may be shown thus : — 

NH 4 OH + C0 2 = NH 4 HC0 3 , 
NaCl + NH 4 HCO s = NaHCO s + NH 4 C1. 



THE ALKALI METALS 211 

The sodium bicarbonate crystallizes out much more quickly than 
the ammonium chloride, and in this way the two compounds are 
separated. The bicarbonate of soda is then heated to expel a portion 
of the carbon dioxide, and sodium carbonate results, thus : — 

2 NaHC0 3 + heat = Xa 2 C0 3 + CO a + H 2 0. 

This process is very cheap because the salt can be had for a few cents 
per hundred pounds, the ammonia is obtained abundantly from all 
gas factories, and the carbon dioxide can be had by calcining limestone 
in making lime. Or, as seen by the last reaction above, the carbon 
dioxide driven off from the bicarbonate of soda may be utilized for 
this purpose, and from the ammonium chloride obtained in the second 
step ammonia may be evolved by treating it with lime. It will be 
seen, therefore, that the result of one part of the process may serve in 
another part and thus reduce the final cost of manufacture. 

The Leblanc Process. — This is more complicated than Solvay's, and 
more expensive ; hence, were it not for the value of some by-products 
which are obtained, it would no longer be used. It really consists of 
three steps. First, common salt is treated with sulphuric acid and 
heated, at first moderately and then more strongly. In the beginning 
the salt is converted into acid sodium sulphate, thus : — 

XaCl + H 2 S0 4 = XaHS0 4 + HC1. 

Next, this acid salt reacts with another part of sodium chloride, form- 
ing normal sodium sulphate, thus : — 

XaCl + NaHS0 4 = Xa 2 S0 4 + HC1; 

or, putting the two together, we have — 

2 NaCl + H 2 S0 4 = Xa 2 S0 4 + 2 HC1. 

The sodium sulphate thus obtained is called salt cake. The hydro- 
chloric acid vapors are passed into flues, down which water constantly 
trickles and absorbs the acid. This is a valuable by-product, and 
serves in some places to keep alive the Leblanc industry. 

Second, this salt cake is mixed with powdered coal and limestone, 
and heated, when sodium carbonate, mixed with several other sub- 
stances, is obtained. The mixture is black in color and is known 
as black ash. The reaction shows the chemical changes that take 
place : — 

Xa 2 S0 4 + CaC0 3 + 2 C = Xa 2 C0 3 + CaS + 2 C0 2 . 



212 MODERN CHEMISTRY 

This black ash is treated with water to dissolve out the sodium car- 
bonate, and the solution is concentrated and purified by calcining, 
dissolving, and recrystallizing. 

In connection with almost every Leblanc factory is also one for the 
manufacture of bleaching powder on a large scale, by using the hydro- 
chloric acid obtained as a by-product, with native manganese dioxide. 

9. Sodium Nitrate, NaN0 3 . — This is known as Chile 
saltpeter on account of the locality from which it is ob- 
tained and its close resemblance to potassium nitrate. 
The crude salt found native in Chile is dissolved in 
water and concentrated, whereupon the pure crystals 
separate. From the fact that it absorbs moisture from 
the air, it cannot be used in making gunpowder to any 
great extent. It is used largely, however, in the manu- 
facture of nitric acid and also for artificial fertilizers. 

10. Sodium Sulphate, Na 2 S0 4 . — This is frequently called 
Glauber s salt. It is a white crystalline salt obtained in 
the preparation of sodium carbonate as described above. 

11. Sodium Bicarbonate, NaHC0 3 . — This is common 
cooking soda, and is usually prepared by , the Solvay 
process of making soda crystals, hence is very inexpen- 
sive. In making bread the " soda " is put with some such 
acid as sour milk. The acetic or lactic acid, or whatever 
it may be, reacts with the soda, setting free carbon dioxide, 
which raises the dough by struggling to escape through it. 
At the same time the acid disappears in the formation of a 
neutral salt. This may be seen by the following reaction 
of " soda " with acetic acid : — 

NaHC0 3 + HC 2 H 3 2 = NaC 2 H 3 2 + H 2 + C0 2 . 

12. Soap. — This is a substance which has been made in 
greater or less quantities for probably two thousand years. 
At first, however, it was used simply as an ointment in a 



THE ALKALI METALS 213 

medicinal way, and not till about 200 a.d. was it applied 
as it is to-day, and even then only to a limited extent. 
Soap is made by combining some alkali, as caustic soda or 
potash, with some fatty substance or oil. The fat con- 
tains an acid which combines with the alkali, hence we see 
that soap is really a salt. It retains some alkaline proper- 
ties, however, just as many other salts do, simply because 
both sodium and potassium are strong alkalies, while the 
fatty acids are comparatively weak ; hence the alkaline 
properties really overbalance the acid properties. It has 
been said that soda crystals, Na 2 C0 3 , are used in making 
soap. They must first be converted into caustic soda, 
however, and this is done by treating the solution with 
milk of lime, Ca(OH) 2 , as described. 

13. Hard and Soft Soap. — We have two kinds of soap, 
hard and soft; the former is made from sodium com- 
pounds, the latter from potassium. Wood ashes contain 
considerable quantities of potassium carbonate ; formerly, 
these were saved by farmers, placed in large " hoppers," 
lime added, and then leached. A dark-colored, strongly 
alkaline solution filtered out, containing a considerable per- 
centage of caustic potash. This was treated with waste 
fat, and boiled, when in the course of a few hours a 
strongly alkaline soft soap was obtained, which always 
remained pasty. By adding common salt to this, it could 
be converted into a dark-colored solid mass ; for many 
years this was the only hard soap known. Sodium com- 
pounds yield hard soap directly on combination with 
fats, hence they are most used at the present time in the 
manufacture of ordinary hard soap. 

14. The practical value of soap lies in the fact that on 
account of its slightly alkaline properties it has the power 
of uniting with the oil secreted by the glands of the skin, 



214 MODERN CHEMISTRY 

and which holds the particles of foreign matter ; this 
" dirt," therefore, may be removed by the mechanical 
action of the water. This also explains why frequent 
bathing with the application of strong soap will tend to 
cause the skin to chap, by the removal of the oil which 
keeps it soft and pliable. 

15. Test for Sodium. — Sodium may always be detected 
by what is known as the flame test. 

Experiment 130. — Heat a platinum wire in the Bunsen flame 
until it no longer imparts any color to the flame. Then dip it into 
the sodium solution and again hold in the flame. The bright yellow 
color is distinctive. 

Potassium : K = 39 

16. Where found. — Because of its great affinity for 
other substances, potassium never occurs free. It is very 
widely distributed, however, in the form of compounds ; 
it is a constituent of many rocks, and by their decomposi- 
tion becomes a part of various soils. Being stored up by 
plants it enters into the animal economy, and by some ani- 
mals, especially sheep, it is exuded from the skin and col- 
lects in considerable quantities upon the wool in an oily 
substance called suint. 

17. How obtained in Metallic Form. — Potassium, like 
sodium, may be obtained by electrolysis, but is usually 
reduced by treating caustic potash with charcoal. The 
reaction shows the chemical changes : — 

6 KOH + 2 C = 2 K 2 C0 3 + 3 H 2 + K 2 . 

The potassium distills out and is collected under oil. 

18. Characteristics of Potassium. —It is a metal some- 
what softer than sodium ; has a bright luster and white 
color, but it tarnishes instantly when cut in the air, so 
great is its affinity for oxygen and moisture. At zero it 



THE ALKALI METALS 215 

becomes crystalline in structure, hard, and brittle. When 
thrown upon water it immediately begins to decompose 
the water, and with such energy that it is melted and 
the hydrogen given off is ignited, burning with a violet 
color. This is due to the vaporization of a small portion 
of the potassium. As hydrogen is set free from the water 
caustic potash is formed, according to a reaction previously 
seen : — j^q + K = K0H + H 

With the halogens, chlorine and bromine, potassium 
ignites spontaneously^ and with liquid ammonia it forms a 
blue solution. It possesses all the strong alkaline charac- 
teristics of sodium in a degree even more marked. 

19. In the metallic form potassium has no uses in the 
arts. Its compounds, however, are very valuable. Any 
potassium salt may be tested in the same way as are the 
sodium compounds, with a platinum wire. The violet 
flame is characteristic. If both sodium and potassium are 
present it will be necessary to observe the flame through 
a blue glass. This transmits the potassium rays, but 
absorbs those of the sodium. 

Compounds of Potassium 

20. Potassium Hydroxide or Hydrate, KOH. — In earlier 
days the most common source of potassium compounds 
was wood ashes, which were boiled with water in iron 
pots. The potassium salts were dissolved out in this man- 
ner, and from them was prepared caustic potash, KOH. 
This is now obtained by a method similar to that used in 
the preparation of caustic soda, viz. by treating potassium 
carbonate with milk of lime, Ca(OH) 2 , when this reaction 
takes place : — 

K 2 C0 3 + Ca(OH) 2 = 2 KOH + CaC0 3 . 



216 MODERN CHEMISTRY 

The latter compound is insoluble and is precipitated ; the 
former is drawn off, concentrated, purified by redissolving 
in alcohol, again dried, fused and molded in the familiar 
round sticks. It is very deliquescent, and quickly dis- 
solves in the moisture it obtains from the air. It is used 
largely as a reagent in the laboratory. 

21. Potassium Carbonate, K 2 C0 3 . — As already indicated, 
this was at one time obtained almost exclusively from the 
ashes of wood. These were treated with water, by which 
the potassium carbonate was dissolved out ; the solution 
was boiled dry, forming a white salt known as pearl ash. 
Now large quantities are obtained by washing sheep's wool 
in hot water, then drawing off the greasy products obtained 
and heating them very strongly to expel the oil. The 
potash salts remain and are dissolved out by water. 

Another source of considerable quantities is the beet- 
sugar industry. The beet sap is boiled down to a sirup, 
and from this sirup is extracted the sugar, leaving a sort 
of molasses, in which still remain the potassium compounds 
that the beets had obtained from the soil. This is gen- 
erally first fermented and distilled; the residue is boiled to 
dryness and calcined. Then from the ashes the potash 
salts are obtained by lixiviation. 

Potassium carbonate may be prepared from the chloride 
by the Leblanc process. 

22. Potassium Chlorate, KC10 3 . — This is a white, crys- 
talline salt, often sold under the misleading name potash. 
It has a not unpleasant, cooling taste, and is used some- 
what for throat affections. In the laboratory it has num- 
berless applications, many of which are familiar to the 
student. In the arts it is used in making matches, for 
fireworks, etc. It is prepared by passing a current of 
chlorine into a solution of caustic potash, by which both 



THE ALKALI METALS 217 

potassium chloride and potassium chlorate are formed. 
The former is much more soluble, hence in concentrating 
the solution the potassium chlorate will crystallize out 
first, leaving the chloride still in solution. 

23. Potassium Nitrate, KN0 3 . — This is commonly known 
as saltpeter. It is a white, crystalline salt, found native 
in various parts of the world. As we have seen, it is pro- 
duced by the decomposition of organic matter, especially 
the refuse from stables. This decomposition is supposed 
to be brought about by the presence of certain bacteria, 
and in some countries the process is now carried on artifi- 
cially to a considerable extent. 

24. The refuse is mixed with ashes and lime, and fre- 
quently stirred to increase the rapidity of decomposition. 
After a time the whole is leached with water to dissolve 
out the nitrate. The solution thus obtained is concentrated 
and the salt allowed to crystallize. 

Considerable quantities are now made by treating sodium 
nitrate, which occurs in almost inexhaustible quantities in 
Chile, with potassium chloride, whereby this double reac- 
tion takes place : — 

KC1 + NaN0 3 = KN0 3 + NaCl. 

Potassium nitrate is used extensively in making gun- 
powder. 

25. Potassium Iodide, KI. — This is a white crystalline 
salt. It is used frequently in the laboratory as a reagent, 
and to some extent in medicine. 

26. Potassium Bromide, KBr. — This is a white salt, 
very similar in general appearance to the iodide. It is 
used frequently in medicine as a sedative. 

Experiment 131. — Take any potassium solution and make the 
flame test just as you did for sodium in Experiment 130. Notice 



218 MODERN CHEMISTRY 

color of the flame. Now mix with it a solution of some sodium com- 
pound, and again test. Can you see the potassium flame? Next 
observe the flame through a sheet of blue glass. State results. 

i 
COMPARATIVE REVIEW OF THE ALKALI METALS 

Sodium and Potassium. 

As found in nature — Two important native compounds of each — 
Where found. 
Which the more important. 
Comparison of the two metals. 
Color. 

Tendency to oxidize. 
Hardness. 
Affinity for water. 
Melting point. 
Affinity for the halogens. 
Experiments that illustrate most of these properties. 

Proof that hydrogen is set free from water by these metals. 
Proof of the hydroxide formed. 
Compounds. 

The Hydroxides — Method of preparing — Reactions. 

Usual form — Appearance — Properties — Uses. 
The Carbonates — Source of supply. 

Former method of obtaining K 2 C0 3 . 
Present sources. 

Two methods of preparing Na 2 C0 3 . 
Uses of the carbonates. 
Review work in glass. 
Kinds of glass — Differences. 
Soap making — Chemistry of. 
Kinds of soap. 
Common salt. 

Preparation for market. 
Cooking soda — Chemical name and formula. 

Chemistry of in bread making. 
Saltpeter — Chemical name and Formula. 

Preparation — Appearance — Uses, 
Potassium chlorate — Formula, 
Appearance — Uses. 



CHAPTER XIX 

THE ALKALINE EARTHS 
Magnesium : Mg = 24 

1. Occurrence. — Magnesium in the form of certain com- 
pounds is widely distributed. Among the most important 
of its compounds may be named the familiar minerals, 
asbestos and meerschaum. The first is a silicate of magne- 
sium and aluminum, and the second a silicate of magnesium. 
Magnesium limestone, or dolomite, CaMg(C0 3 ) 2 , occurs in 
considerable quantities. 

2. Peculiarities of the Metal. — Magnesium is silvery 
white in color, and melts at a red heat. In dry air it does 
not tarnish, but moisture quickly affects it. While at 
ordinary temperatures it is slightly brittle, as it nears the 
melting point it becomes malleable and may be drawn into 
wires. These, flattened into ribbons, are the usual com- 
mercial form, though the powder is also frequently seen. 
The metal is easily ignited and burns with a dazzling 
white light, rich in actinic properties. This combustion 
is so vigorous that it will decompose even carbon dioxide 
and certain other similar oxides. (See carbon dioxide, 
page 143.) 

3. Uses. — On account of the light furnished by burn- 
ing magnesium, it is frequently used in taking flash-light 
pictures of caverns, and other interior views. It is like- 
wise used to a limited extent in making fireworks. In 
the form of a powder it is often used like zinc in the 

219 



220 MODERN CHEMISTRY 

reduction of ferric to ferrous salts (see page 307), on 
account of the rapidity with which, in the presence of sul- 
phuric or hydrochloric acid, it yields hydrogen. It is 
also used by chemists in cases of supposed arsenic poison- 
ing, in making Marsh's test. Zinc nearly always contains 
traces of arsenic, whereas magnesium is obtained prac- 
tically pure ; for this reason it is substituted for the zinc. 

Compounds of Magnesium 

4. Magnesium Sulphate, MgS0 4 . — One of the most 
common compounds is epsom salts, magnesium sulphate, 
MgS0 4 , 7 H 2 0. This is a salt found in the water of many 
mineral springs. It has a very bitter taste and is used 
largely in medicine, also extensively in finishing cotton 
goods. 

5. Magnesia, Magnesium Oxide, MgO. — This is a white 
solid obtained when magnesium is burned in the air or in 
oxygen. It is often prepared by heating magnesium car- 
bonate to expel the carbon dioxide, just as lime is pre- 
pared from limestone (see lime, page 221). The reaction 
is seen below : — 

MgC0 3 + heat = MgO + C0 2 . 

It is used as a face powder, and, because of its high melt- 
ing point, sometimes for making or lining crucibles. 

Calcium: Ca = 40 

6. Occurrence. — In the form of compounds, calcium is 
one of the most abundant and most widely distributed 
elements known. Because of its strong affinity for water, 
however, it never occurs free. The carbonate of calcium, 
CaCOg, is the most abundant form and includes many well- 



THE ALKALINE EARTHS 221 

known substances, such as marble, limestone, and chalk. 
Some of the more highly crystallized forms are Iceland 
spar, calcite, and dog-tooth spar, while stalactites, corals, 
and shells have the same composition. The next most 
abundant natural compound of calcium is gypsum, calcium 
sulphate, CaS0 4 , 2 H 2 0. 

7. Production of the Metal. — Calcium has seldom been 
prepared, and then only for the purpose of studying its 
properties. Sir Humphry Davy, who first isolated potas- 
sium and sodium from their hydroxides by means of an 
electric current, in the same way decomposed calcium 
chloride and obtained calcium in the metallic form. 

8. Characteristics. — Calcium is of a brassy yellow color, 
and somewhat malleable and ductile. It has a density of 
about 1.6, and like sodium readily decomposes water, 
forming the hydroxide, Ca(OH) 2 . It is readily soluble in 
dilute acids, and at a temperature a little above its melt- 
ing point it burns with a reddish yellow light. The cost 
of its production, 15.00 a pound, precludes its practical 
use. 

Compounds of Calcium 

9. Although as a metal calcium is of so little value, it 
would be difficult to estimate the worth of the compounds. 

10. Lime, Calcium Oxide, CaO. — This is one of the 
most important compounds known. It is easily prepared 
from limestone by heating it to a red heat, at which tem- 
perature carbon dioxide is expelled, thus : — 

CaC0 3 + heat = CaO + C0 2 . 

Lime is prepared in kilns, which are simply square rooms 
or ovens 15 to 20 feet high, and 10 to 15 feet each way. 
See Fig. 52. The limestone is thrown in from above and 



222 



MODERN CHEMISTRY 




Fig. 52. 



strongly heated with dry cordwood or coke in alter- 
nate layers. In a few hours the limestone is converted 

into lime, then the fire is 
removed, the mass is allowed 
to cool and the lime with- 
drawn, and if intended for 
shipment packed in barrels. 
Some kilns are arranged 
below so as to enable the 
workmen to remove the 
lime without putting out 
the fire. Such are contin- 
uously fed from above, and 
the operation goes on with- 
out ceasing. 

11. Properties of Lime. — Prepared as above it is in the 
form of white lumps, but if left exposed to the air it 
begins at once to take up moisture and in a short time 
crumbles to a fine powder. It is then said to be " air- 
slaked," although it is really the water in the air that has 
caused the change. The reaction is as follows : — 

CaO + H 2 = Ca(OH) 2 . 

12. If a lump of freshly prepared lime be treated with 
water, the change indicated above takes place rapidly, 
accompanied by the evolution of considerable heat. The 
hydroxide, Ca(OH) 2 , thus obtained is soluble in water, 
though very much less so than ordinary caustic soda or 
potash. The solution of caustic lime is known as lime- 
water. 

13. Uses of Lime. — Lime is indispensable in the erec- 
tion of almost all structures. Mixed with sand it forms 
the mortar for nearly all stone and brick work — except 



THE ALKALINE EARTHS 223 

such as is laid under water — and much of the plaster for 
indoor work. Unmixed with sand it is frequently used 
to give the white or finishing coat in plastering, though 
various plasters are now beginning to take the place of 
ordinary lime in this respect. 

14. It is also used extensively in the lime purifiers of 
illuminating gas works, in the manufacture of bleaching 
powder, of ammonia, in removing the hair from hides in 
the process of tanning, and for numerous other purposes 
where a cheap and easily prepared alkali is demanded. 

15. J^alcium liydroxide, exposed to the air, absorbs car- 
bon dioxide and forms calcium carbonate, thus : — 

Ca(OH) 2 + C0 2 = CaC0 3 + H 2 0. 

The same reaction takes place in mortar, hence that which 
has been properly prepared should grow gradually harder, 
in time being converted back again into a siliceous lime- 
stone. If a beaker containing lime-water be left exposed 
to the air, in a little while a white film will be seen to 
cover the surface of the liquid. This is really a pre- 
cipitate of calcium carbonate, resulting from the absorp- 
tion of the carbon dioxide of the air by the lime-water. 
If the breath from the lungs be blown through a clear 
solution of lime-water, it quickly becomes clouded from 
the same cause. 

16. Calcium Carbonate, CaC0 3 . — In the natural form 
this is known in the several varieties mentioned above. 
Artificially, it may be prepared as a white precipitate by 
adding some alkaline carbonate, as sodium or ammonium 
carbonate, to a solution of calcium. The following reac- 
tion takes place : — 

CaCl 2 + (NH 4 ) 2 C0 8 = 2 NH 4 C1 + CaC0 3 . 



224 MODERN CHEMISTRY 

It is insoluble in pure water, but when an excess of carbon 
dioxide is present, it slowly dissolves. 

Experiment 132. — Through a few cubic centimeters of lime-water 
in a flask or beaker, pass a current of carbon dioxide, or blow the 
breath for some time. What finally becomes of the white precipitate 
which forms at first ? Preserve the water. 

In this way water charged with carbon dioxide percolating through 
limestone rocks gradually dissolves them, and has formed many of the 
great caves known in this country. This same water, dripping from 
the roof of caverns, being no longer under pressure, gives up its carbon 
dioxide, and the calcium carbonate, no longer held in solution by the 
gas, is deposited in the form of stalactites and stalagmites. 

17. Calcium Chloride, CaCl 2 . — This is a white salt which 
may be prepared from any form of the carbonate by treat- 
ing with hydrochloric acid. It is a by-product formed in 
the preparation of carbon dioxide from limestone : — 

CaC0 3 + 2 HC1 = CaCl 2 + C0 2 + H 2 0. 

It is strongly deliquescent, and is often used in drying 
gases, damp cellars, etc. 

18. Calcium Sulphate, CaS0 4 , 2 H 2 0. — In the natural 
form this is the gypsum already mentioned. It occurs 
in vast quantities in many of our states, notably Kansas, 
New York, Illinois, etc., both in the form of rich, heavy 
deposits, and mixed with various impurities upon the 
surface. It is used extensively in making plaster of 
Paris. This is manufactured simply by strongly calcining 
the powdered gypsum until the water of crystallization is 
expelled. During this time, as the water escapes from 
the powdered mass, the whole seems to boil vigorously. 
After two or three hours the process is complete, and the 
plaster is ready to be mixed with the " retarder," if neces- 
sary. This plaster has the property of " setting " or hard- 



THE ALKALINE EARTHS 225 

ening quickly when water is added to it. This is due to 
the fact that the anhydrous salt again takes up the water 
of crystallization expelled in the previous calcination. If 
the plaster which has been used once be again calcined, 
it acquires again its property of "setting." 

19. Uses of Plaster of Paris. — It is employed extensively 
in making molds for many of the finer castings, in dental 
work and surgery, for statuettes, as a finishing coat in 
plastering, and for stucco and other ornamental work on 
the interior of buildings. For most purposes, a plaster 
that does not harden so rapidly is desirable, hence it is 
customary to mix with it some kind of clay, or other 
substance, which causes it to " set " more slowly. This 
clay has already been spoken of as the "retarder." 

20. Cements. — Cements are a species of lime which 
have the power of hardening or setting rapidly, like 
plaster of Paris. They are prepared by calcining lime- 
stone, which contains a large percentage of silica and 
alumina, Si0 2 and A1 2 3 . Dolomitic or magnesium lime- 
stones, containing also the silica and alumina, when cal- 
cined, produce a cement that will harden under water, 
known as hydraulic cement. It has been stated that 
ordinary plaster hardens by the absorption of carbon diox- 
ide from the air, forming again calcium carbonate. This 
is, necessarily, a slow process. Cements,, as already stated, 
are produced by driving out the water of crystallization ; 
hence, when they are mixed with water for use, they very 
rapidly take this up again, forming practically the original 
rock. Hydraulic cement is used in laying the piers of 
bridges, building jetties, and other work that is to be 
under water. Ordinary cements are used extensively for 
laying pavements, building roadbeds, for the concrete 
foundation for various kinds of masonry, etc. The fol- 



226 



MODERN CHEMISTRY 



lowing shows the composition of some cement rocks from 
various localities : — 



Locality 


CaC0 3 


MgC0 8 


Si0 2 


Fe 2 3 


A1 2 3 


H 2 


Undeter- 
mined 


Rosendale, N.Y. 


45.91 


26.14 


15.37 




11.38 


1.20 




Utica, 111. 


42.25 


31.98 


21.12 




1.12 


1.07 


2.46 


Milwaukee, Wis. 


45.54 


32.46 


17.56 


3.03 


1.41 






Cement, Ga. 


43.50 


22.00 


22.10 


1.80 


5.45 


4.95 




Siegfried, Pa. 


78.90 


2.66 


11.62 




6.25 




0.55 


Ft. Scott, Kan. 


73.95 


2.26 


18.75 


2.32 


2.15 


0.37 


0.20 


Ft. Scott, Kan., No. 2 


65.21 


10.65 


15.21 




4.56 




4.37 



21. Hard Water. — Hardness in water is due to the 
presence of certain salts in solution, very commonly some 
compounds of calcium. This hardness may be either tem- 
porary or permanent, according as it may be removed by 
boiling or by adding ammonia, or not at all. 

Experiment 133. — Prepare a soap solution by dissolving a shaving 
of soap in warm distilled water. Allow it to stand a few minutes. It 
should be perfectly clear. To a few cubic centimeters of the lime- 
water, through which the breath was blown till clear again, add a 
little of the soap solution. What happens? Why? Take another 
portion of the clear lime-water and boil it for a few minutes. Has 
any sediment formed in the flask ? The heat has expelled the carbon 
dioxide ; why does the precipitate form ? Decant a portion of it and 
test with the soap solution : is the water still " hard " ? What effect 
has the boiling had ? 

To another portion of the same hard water (which has not been 
boiled) add a few drops of ammonia and again test to see whether the 
water is still hard. What are the results? 

Add a little powdered calcium sulphate, CaS0 4 , to some water, and 
after some time test a portion of it to learn whether it is hard. 
Now try to remove the hardness by the methods previously used. 
State results. 



TEE ALKALINE EABTHS 227 

22. Water the hardness of which may be easily re- 
moved is said to be temporarily hard, while that which 
cannot be so changed is permanently hard. When the 
hands are washed with soap in hard water, the soap pre- 
cipitates the salts in the water, of which a portion settles 
upon the skin, giving it an unpleasant feeling. Another 
part of the precipitate is usually seen as a scum upon the 
surface of the water. 

23. Bleaching Powder, Ca(C10) 2 + CaCl 2 . — This is also 
called hypochlorite of lime. It is a white powder which 
is prepared by passing chlorine into chambers containing 
common lime spread loosely upon shelves. The reaction 
may be represented thus : — 

2 CaO + 4 CI = Ca(CfO) 2 + CaCl 2 . 

When treated with any dilute acid, chlorine is again set 
free ; for this reason the compound is used extensively as 
a source of chlorine in bleaching muslin and other cotton 
goods.* 

24. From the fact that chlorine does not bleach dry 
cloth, it is believed to be not the direct bleaching agent, 
but simply that which sets free another. It will be seen 
later, in studying the compounds of manganese, that log- 
wood, litmus, and other colored vegetable solutions are 
rapidly bleached by the use of potassium permanganate, 
in the presence of some acid. Experiment shows that 
this is due to the oxygen that is set free from the per- 
manganate. Similarly the chlorine, which has most won- 
derful affinity for hydrogen (see page 108), sets free the 
oxygen from the water with which the cloth is moistened, 
and this in the nascent state oxidizes the coloring matter 
and reduces it to colorless compounds. 

* See work under Chlorine, page 111. 



228 MODERN CHEMISTRY 

25. When a current of carbon dioxide is passed through 
a solution of bleaching powder, chlorine is liberated, and 
can be detected by the odor, just as when treated as above 
with a dilute acid. Exposed to the air, bleaching powder 
yields up its chlorine, owing to the action of the carbon 
dioxide always present ; but naturally the process is very 
slow. On account of this fact, and because chlorine is an 
excellent germicide and disinfectant, bleaching powder is 
used frequently in sick rooms and hospital wards. The 
generation of the chlorine is so slow as to be scarcely 
noticeable, and yet sufficient to keep the atmosphere in 
a wholesome condition. 

Strontium : Sr = 87 

26. Its Name. — Strontium is a rare metal, which re- 
ceived its name from Strontian, a place in Scotland, where 
it was discovered. One of its chief sources is the mineral 
strontianite, SrC0 3 . 

Compounds of Strontium 

27. Strontium Nitrate, Sr(N0 3 ) 2 . — This is a white crys- 
talline salt, soluble in water. It is used considerably in 
fireworks and in making " red fire." 

Experiment 134. — Mix thoroughly about a gram each of stron- 
tium nitrate and potassium chlorate, finely pulverized, and about as 
much in bulk of powdered shellac. Place the mixture in an iron 
saucer and ignite with a match. State the results. 

28. Strontium Carbonate, SrC0 3 . — This is a white pre- 
cipitate, like calcinm carbonate, obtained when ammonium 
or sodium carbonate is added to a neutral or alkaline 
solution of a strontium salt. 

Sr(N0 3 ) 2 + (NH 4 ) 2 C0 3 = SrC0 3 + 2 NH 4 N0 3 . 



THE ALKALINE EARTHS 229 

29. Strontium Hydroxide, Sr(OH) 2 . — When water is 
added to strontium oxide, SrO, like lime, it is slaked, 
evolves much heat, and is converted into the hydroxide, 
Sr(OH) 2 . In this form it is used considerably in the 
manufacture and refining of beet sugar. 

Barium : Ba = 137 

30. Its Name. — This metal, also rare, received its name 
from a Greek word, meaning heavy, and was so called be- 
cause its chief natural ore, baryta, BaS0 4 , has great den- 
sity. It is also found as a carbonate, BaC0 3 , known as 

wither ite. 

Compounds of Barium 

31. Barium Chloride, BaCl 2 . — This is a white crystal- 
line salt, readily Soluble in water. It is used in the 
laboratory in te^titig for sulphuric acid. 

32. Barium Sulphate, BaS0 4 . — This is a heavy white 
precipitate, insoluble in water and acids. It is easily pre- 
pared by adding sulphuric acid or any soluble sulphate to 
a solution of barium chloride. It is used considerably as 
an adulterant for white lead (see page 280), and to some 
extent in weighting paper. 

33. Barium Carbonate, BaC0 3 . — This is a white precipi- 
tate formed when ammonium or sodium carbonate is added 
to a neutral or alkaline solution of a barium salt. It is 
insoluble in water, but soluble in weak acids. 

Experiment 135. — Let the student prepare both of the above 
compounds, using barium chloride for the barium solution. Note the 
differences between the two and test their solubility in the common 
acids. State results. 

34. Barium Nitrate, BaN0 3 . — This is a white crystal- 
line salt. It is used to a considerable extent in the 
making of green fire for fireworks. 



230 MODERN CHEMISTRY 

Experiment 136. — Repeat Experiment 134, substituting barium 
nitrate for the strontium nitrate, and state results. Sulphur or pow- 
dered charcoal may be used instead of the shellac, but the sulphur 
yields very irritating fumes of the dioxide, and the charcoal does not 
burn so readily. 

35. Barium Hydroxide, Ba(OH) 2 . — This is a compound 
obtained from barium oxide, BaO, by the addition of 
water, just as slaked lime is prepared. Like calcium 
hydroxide, it forms a precipitate of the carbonate upon 
the addition of carbon dioxide. It was formerly used 
extensively in clarifying beet sugar, but as it is very 
poisonous, and traces of it sometimes remain in the sugar, 
its use has been supplanted by that of strontium hydroxide. 

36. Flame Tests. — The metals of this group, calcium, 
strontium, and barium, may be detected by the flame test. 

Experiment 137. — Just as you tried sodium and potassium, how 
take some solutions of these three' metals and make the flame test in 
the same way. State results as to color and duration of flame. 

REVIEW OF WORK IN ALKALINE EARTH METALS 

Magnesium, Calcium, Strontium, Barium. 

1. Occurrence — Compare native compounds. 

Crystallized forms of calcium compounds. 
Un crystallized forms. 
Special forms. 

2. Artificial compounds. 

a. The Oxides of Mg, Ca, Sr, Ba. 

Wherein is their preparation similar? Why are 

such compounds used ? 
Importance of CaO and MgO. 

b. The Hydroxides — Similarity of preparation. 

Uses of Ca(OH) 2 and Sr(OH) 2 . 

Preparation of mortar ; chemical change it under- 
goes as it hardens. 
Hydraulic cement ; other cements ; uses ; explanation. 



THE ALKALINE EARTHS 231 

c. The Nitrates — Use in the arts of Sr(N0 3 ) 2 ; Ba(X0 3 ) 2 . 

How used. 
Chemical action of each constituent. 

d. The Sulphates — Two important ones. 

Preparation of Plaster of Paris — Compare with prepa- 
ration of CaO. 

Uses of CaS0 4 and BaS0 4 . 

Chemical change which takes place when Plaster of 
Paris hardens. 

Compare with hardening of mortar. 

e. Hard Waters — Due to what compounds. 

Two classes, how different. 
Methods of softening water. 
Chemistry of these methods. 
/. Some special calcium compounds. 

CaF 2 — Use, and method of using. 

Bleaching powder — Uses — Compare CI and S0 2 as 
bleaching agents — Chemical action of each. 

Use of bleaching powder as a disinfectant — 
How is chlorine set free? 

3. Flame tests — Method of making test. 

Comparison of colors imparted. 

4. Comparative value of the metals in metallic form. 

5. Exercise — Given some marble, HC1, H 2 S0 4 , H 2 0, and Xa 2 C0 3 . 

Tell how to prepare CaO, Ca(OH)" 2 , CaS0 4 , CaCl 2 , C~aC0 3 
(amorphous). Write all reactions concerned. 



CHAPTER XX 

THE COPPER-SILVER GROUP — COPPER, SILVER, GOLD 

Copper : Cu = 63 

1. History. — Copper has been known from earliest 
antiquity, its use being mentioned by Jewish, Assyrian, 
and other ancient historians. By the Greeks it was ob- 
tained from the island of Cyprus, and from this fact 
probably received the name kuprum, and its present symbol, 
Cu. In England copper-mining was begun before the close 
of the twelfth century. It met with little success, however, 
till about five hundred years later. In the United States, 
the oldest mines are those of the Lake Superior region. 
The remains of prehistoric tribes about the mines indicate 
clearly that these deposits were known and used in very 
early times. The metal was obtained by stripping the 
rock and earth from the outcropping strata. When the 
rock had been broken or cracked off, the thin sheets of 
copper were removed and hammered into vessels of various 
shapes. 

2. Sources of Supply. — Besides the mines of northern 
Michigan, which yield almost pure copper, large quantities 
are obtained from the silver ores of Montana and Colorado. 
Many of the mines of Michigan are exceedingly productive, 
some of them yielding annually about 25,000 tons, but in 
recent years the mines of Montana have furnished about 40 
per cent of the world's supply. Among the ores found in 
the Western mines may be mentioned malachite, CuC0 3 , 
Cu(OH) 2 , green in color; azurite, CuC0 3 , 2 Cu(OH) 2 , 

232 



THE COPPER-SILVER GROUP 233 

a beautiful blue, usually associated with the malachite; 
chalcopyinte, or copper pyrite, CuFeS 2 , a brass-colored 
ore, resembling fool's gold, but often having a purplish 
cast ; and bomite, a sulphide of iron and copper of varying 
proportions. 

3. Reduction of the Ore. — In the case of the copper 
from the Lake Superior mines, scarcely any refining is 
necessary. It is passed through crushers to break up the 
rock associated with the metal, then by washing and other 
mechanical processes the separation is effected. 

4. Methods in the West. — When the ore is a carbonate, 
like malachite or azurite, or the oxide, it is simply mixed 
with coke and reduced according to the general plan of 
reducing metallic ores. Thus, 

CuO + C = Cu + CO. 

Usually, however, there is a high per cent of sulphur 
present, and the process is much more complicated. There 
are, in reality, four stages necessary before blister copper, 
that is copper about 98.8 per cent pure, is obtained. These 
four are concentration, calcination, reverberation or blast re- 
duction, and converting. The first consists in the separation 
of the silica or rock from the copper ore. This is done 
by mechanical washing with " jiggers." By calcination 
the sulphur is partially removed. After the ore has been 
roasted, either one of two plans may be followed. Accord- 
ing to one method, the red-hot ore is placed in reverbera- 
tory furnaces and melted. The sulphide of copper, mixed 
with the sulphide of iron, always present, and the silver 
and gold, being heavy, settle to the bottom. This molten 
mixture is drawn off and is known as matte. 

5. Sometimes the ore, even without concentration or 
calcining, is put directly into blast furnaces. In this case 



234 MODERN CHEMISTRY 

limestone rock is mixed with the ore ; when the mass is 
heated the silica and limestone unite to form a glassy slag 
which takes up about 75 per cent of the iron. The slag, 
being relatively light, is drawn off above the metal. The 
sulphur in excess is removed by the strong draughts 
of air which are forced through the blast furnace. A 
matte is thus obtained similar in composition to that pro- 
duced by reverberation. 

6. The fourth stage consists in converting this matte 
into blister copper. This is done in a converter, which 
in its essentials is not unlike the Bessemer converter 
described in detail in the chapter on iron. The molten 
matte has fine streams of air driven through it, and in 
a few minutes is converted into copper about 98.5 per 
cent pure. This still contains small quantities of iron, 
arsenic, gold, and silver, which are finally separated at 
the refineries. 

Experiment 138. — Put upon charcoal a little copper oxide, CuO, 
mixed with sodium carbonate, and heat strongly in the reducing flame. 
Note the color of the granular mass remaining. Test its malleability 
with a hammer. What have you obtained? 

7. Characteristics of Copper. — Copper is a very tena- 
cious, malleable, ductile metal, of a reddish color. It does 
not tarnish in dry air, but in the presence of moisture and 
carbon dioxide is slowly converted into a green carbonate 
of copper. With the exception of silver it is the best 
conductor of electricity known. Its melting point is high, 
being nearly 1100° C. In the oxidizing flame it is con- 
verted into the black oxide of copper, CuO. It is soluble 
in nitric acid and in hot concentrated sulphuric acid. 
From its solutions it is easily precipitated by iron, zinc, 
and certain other metals. 



THE COPPER-SILVER GROUP 235 

8. Applications in the Arts. — With the exception of 
iron, copper, probably, has more varied uses than any 
other metal. It is employed very extensively in alloys, 
among them being the following : — 

Brass : consisting of copper and zinc in varying pro- 
portions. 

Bronze : copper, zinc, and tin. 

Bell-metal : copper and tin. 

Coinage : gold and silver with copper. 

Aluminum bronze : aluminum and copper. 

A peculiarity of the last is that, with about 1 to 3 
per cent of copper, it is of a beautiful silver-white color, 
much whiter even than aluminum; with 10 per cent of 
copper it somewhat resembles gold. In the latter propor- 
tions it is used largely for making various fancy articles 
and novelties. 

9. Unalloyed, copper is used for roofing, for the sheath- 
ing of vessels, for making various utensils, and for wire for 
trolley, telegraph, and telephone sj'stems, and for electric 

lighting. 

Compounds of Copper 

10. Two Classes of Salts. — Copper, like several other 
metals, forms two classes of salts, cuprous and cupric, 
though as a rule only the latter are of importance. 

11. Cupric Sulphate, CuS0 4 , 5 H 2 0. — This is commonly 
known as Hue vitriol. It forms in beautiful blue crystals, 
and is obtained when metallic copper is dissolved in boil- 
ing sulphuric acid. The commercial supply is obtained 
mostly as a by-product from the great gold and silver 
refineries, such as those of Kansas City and Omaha. The 
smelters at the former place produce monthly about eight- 
een hundred tons, worth between $100,000 and $200,000. 
The silver ores contain more or less copper in the form oi 



236 MODERN CHEMISTRY 

cupric sulphide, which in the roasting of the ore is con- 
verted into cupric sulphate. 

CuS + 2 2 = CuS0 4 . 

This, being soluble in water, is washed out and concentrated, 
whereupon the crystals separate out from the solution. 

12. Characteristics and Uses. — The salt is somewhat 
efflorescent, and when exposed to the air gradually gives 
up a portion of its water of crystallization. At the same 
time it breaks up and becomes almost white in color. By 
heating, the water of crystallization may be entirely re- 
moved and the blue color destroyed. This may be 
restored, however, by digesting for some time in water. 
Blue vitriol is very poisonous, and is used extensively in 
making Paris green and Bordeaux mixture for spraying 
fruit trees to destroy moths and other insects. It is 
employed largely in electroplating and electrotyping, also 
in galvanic batteries, though the dynamo is now taking 
the place of these batteries. 

13. Cupric Nitrate, CuN0 3 , 3 H 2 0. — This is a deep blue 
solid, soluble in water, obtained when copper is treated 
with nitric acid. 

14. Cupric Chloride, CuCl 2 . — This is a beautiful tur- 
quoise-blue, finely crystallized salt. 

15. Cupric Sulphide, CuS. — This is a black precipitate 
obtained when a current of hydrogen sulphide is passed 
through a solution of a copper salt. It is soluble in hot 
nitric acid, and partially so in warm yellow ammonium 
sulphide. 

16. Cupric Acetylide, CuC 2 , H 2 0, or CuC 2 . — Cupric 
acetylide is a reddish brown precipitate formed when 
acetylene is passed through a copper solution. In drying 
it gives up its molecule of water and becomes very explo- 



THE COPPER-SILVER GROUP 237 

sive, a slight jar being sufficient to touch it off. Metallic 
copper which has for some time been in contact with moist 
calcium carbide is partially converted into the acetylide, 
and shows the same explosive tendencies. 

17. Cupric Oxide, CuO. — This is a black powder, ob- 
tained when copper is heated to redness in the air, or 
when cupric nitrate is treated in a similar manner. In 
the hydrated form, CuO, H 2 0, it may be obtained by 
treating a copper solution with caustic soda or potash 
and boiling for a few minutes. 

Experiment 139. — To prepare some of the above compounds. 
The nitrate and sulphate have already been prepared. Review the 
work with nitrogen dioxide and sulphur dioxide. 

Add to a few cubic centimeters of copper nitrate solution a few 
drops of ammonium sulphide, (XH 4 ) 2 S, or pass through it a current 
of hydrogen sulphide. Xote the color of the precipitate formed. What 
is it? 

Put into a crucible or small evaporating dish a half gram of pow- 
dered copper nitrate, and heat gradually to dull redness. How is the 
nitrate changed? What gas did you see expelled? What have you 
obtained? Save the powder. 

Put into a test-tube a few cubic centimeters of a solution of copper 
nitrate or sulphate, and add a little caustic soda or potash. A blue 
precipitate of cupric hydrate is obtained, Cu(OH) 2 . Boil it for a few 
minutes. Notice the change in color. What have you obtained? 

Make a borax bead upon a platinum wire and fuse into it a little 
of the cupric oxide prepared above. ^Y r hat colored bead do you ob- 
tain? The oxide is thus used sometimes in preparing emerald glass. 

Experiment 140. Practical Work. — To determine the composi- 
tion of brass. Dissolve a few brass filings in warm nitric acid. Notice 
the color of the solution obtained. What metal is indicated by the 
color? Evaporate nearly to dryness, and take up with 40 to 50 cc. of 
water. Warm gently and pass a current of hydrogen sulphide for 
several minutes, or until no further precipitate will form. This may 
be determined by filtering out a little and passing the gas through it. 
If no precipitate forms, the whole may be filtered. Punch a hole in 
the bottom of the filter as it rests in the funnel, and wash the black 



238 MODERN CHEMISTRY 

precipitate through into a beaker with a little nitric acid diluted. 
Heat until it dissolves. What is indicated by the color of the solu- 
tion ? To prove, add ammonia until alkaline. Do you obtain a deep 
blue solution? If so, copper is indicated. 

The nitrate obtained above from the black precipitate will con- 
tain the other metal or metals found in the brass. Add to it a 
few drops of ammonium hydroxide and then a little ammonium 
sulphide. Do you obtain a starchy white precipitate ? If so, zinc is 
indicated. 

Exercise. — Write reactions showing the preparation of cupric 
sulphate, nitrate, sulphide, acetylide, hydrate, oxide, and the reactions 
in the analysis of brass as far as possible. 

Silver : Ag = 108 

18. Ores of Silver. — This metal has been known from 
remote antiquity, because of the fact that it frequently 
occurs free in small particles disseminated through quartz 
and other rock. Occasionally large masses have been 
found, and in the museum at Copenhagen there is to be 
seen one weighing about five hundred pounds. Usually, 
however, silver is in combination with other elements. 
One of the most important ores is horn silver, AgCl, named 
from its general resemblance in color and texture to the 
horns of cattle. Another important ore is argentite, Ag 2 S. 
As the greater part of the lead ore smelted contains more 
or less silver, lead furnaces yield the largest portion of 
the world's output of silver. 

19. Reduction of the Ores. — The following experiment 
will illustrate roughly one of the methods by which silver 
ores are reduced. 

Experiment 141. — To about 10 cc. of a solution of silver nitrate, 
add a little hydrochloric acid. The precipitate is silver chloride, 
AgCl ; shake the contents, warm slightly, and when the precipitate 
has settled, decant the moderately clear solution. Transfer the curdy 
white precipitate to a piece of charcoal, cover with sodium carbonate, 



THE COPPER-SILVER GROUP 239 

and heat strongly in the reducing flame. Presently a bright globule 
of silver will appear. This may be preserved for tests upon the 
metal if desired, or dissolved in dilute nitric acid and converted again 
into silver nitrate. 

20. Other Methods. — Various processes are used in 
reducing silver ores, depending upon the character of the 
ore. But so. large a proportion of the silver output results 
from lead reduction, that we shall confine ourselves here 
to only one or two of the methods employed. When these 
argentiferous lead ores are reduced (see Lead, page 275), 
the two metals, silver and lead, are formed together as an 
alloy, and they must then be separated. There are two 
methods for doing this. When the alloy is rich in silver, 
Pattison's method is employed. 

21. Pattison's Method. — It has been found that when 
such an alloy is allowed to cool slowly the lead will crys- 
tallize before the silver. Hence, as the lead crystals begin 
to form they are skimmed out with perforated ladles, thus 
dividing the alloy into two portions, one containing the 
silver with a very little lead remaining in it, and the other 
the lead, with very small quantities of silver. The first 
of these is then submitted to cupellatioji. The alloy is 
gradually run into a cupel, or cup, which is placed upon a 
hearth within the furnace. A blast of air and flame is 
directed upon the surface of the alloy, and the lead is 
oxidized to litharge, PbO. The current of air constantly 
drives this film of oxide off into another vessel so placed 
as to receive it. In this way the lead is entirely removed, 
and the completion of the process is known by the brilliant 
appearance of the molten silver. 

22. Parke's Process. — Zinc will readily alloy with sil- 
ver but not with lead, and this principle is made use of 
in Parke's process of separating lead and silver. Zinc is 



240 



MODERN CHEMISTRY 



added to the alloy, and the whole is melted. The alloy of 
zinc and silver, being lighter than the lead, rises to the 
surface, and as it begins to solidify is skimmed off in the 
form of crystals. Thus there is obtained an alloy of zinc 
and silver with very little lead adhering. This alloy is 
now very carefully heated in a furnace, the bottom of 
which is inclined ; the lead melts and runs off before the 
fusing point of the alloy is reached. The zinc still 
remaining is next removed by heating strongly in retorts, 
when it is vaporized and passes off. 

Experiment 142. — Making use of the bead of silver obtained 
above in Experiment 141, test its hardness and malleability. Try to 
oxidize it in the oxidizing flame. Does any coating form upon the 
charcoal? For just a moment put a silver coin into a solution of 
hydrogen sulphide or sodium sulphide. What are the results? Next 
immerse it in a moderately strong solution of potassium cyanide, and 
allow it to remain some time, if necessary. State the results. This 
last suggests a method of cleaning tarnished silverware, but it should 
be used with caution, as the cyanide is deadly poison. 

Experiment 143. — Add to 2 or 3 cc. of silver nitrate a little hydro- 
chloric acid, spread the white precipitate smoothly upon a sheet of 

paper, place upon it any figure cut 
from thick paper, and expose it to the 
light. In a few minutes, notice what 
has happened. This illustrates the 
method of printing from photographic 
negatives upon sensitized paper. 

The experiment may be varied, and 
with care and patience most beauti- 
ful prints may be obtained. Immerse in a solution of silver nitrate a 
sheet of drawing paper, and allow it to dry in the dark. Next immerse 
in a solution of common salt, and again let it dry in the dark. When 
ready to print, place upon this paper, thus sensitized, an old negative, 
or even a fern leaf or any similar object, and expose to bright sun- 
light, under a sheet of glass to hold in place. Notice when a deep pur- 
ple is obtained, then immerse in a solution of sodium thiosulphate, the 
photographer's " hypo," and rinse thoroughly in water several times. 




Before Fig. 53. After 
Exposure. Exposure. 



THE COPPER-SILVER GROUP 241 

23. Characteristics of Silver. — Silver is a white, lustrous 
metal, malleable and ductile, an excellent conductor of 
electricity and heat, of medium hardness and density. 
It is quickly attacked by many sulphur compounds and 
by the members of the halogen group, although it does 
not tarnish in the air at any temperature. In living 
rooms silverware is tarnished by the action of the sulphur 
gases thrown off in the combustion of coal or of ordinary 
illuminating gas. Eggs and various other articles of food 
tarnish silverware for a similar reason. What is known 
as " oxidized " silver is really that which has been treated 
with some compound of sulphur, producing silver sulphide 
upon the surface. 

24. Uses for Silver. — Owing to its brilliancy and dura- 
bility, silver has long been used for jewelry and various 
other articles of ornament. Alloyed with some other 
metal to make it harder, it is employed extensively in 
coinage ; is also used in amalgams for dentistry and for 
the backs of high grade mirrors, and for plating innu- 
merable articles of use and ornament. 

Compounds of Silver 

25. There are only a few compounds that are of interest, 
and but one or two that are of any considerable value. 

26. Silver Nitrate, AgN0 3 . — This is important because 
most of the other silver compounds are prepared from it, 
and because it has numerous applications in the arts. It 
occurs in slab-like, almost transparent, white crystals, 
which are soluble in water. It is prepared by dissolving 
silver in nitric acid. When exposed to the light, espe- 
cially if in contact with any organic matter, it turns dark. 
It is used for sensitizing paper for photographic work, as 
the principal ingredient of indelible ink, and in hair dyes. 



242 MODERN CHEMISTRY 

In the form of lunar caustic, which is simply crystallized 
silver nitrate fused and molded into sticks, it is used in 
cauterizing wounds, such as dog bites, for ulcerated sore 
throat, in removing warts and other similar excrescences 
of the skin. 

27. Silver Chloride, AgCl. — This is prepared from a 
solution of silver nitrate by adding to it hydrochloric acid 
or any soluble chloride, like common salt. It is a white 
precipitate, curd-like in appearance, especially when shaken 
for a moment. It is soluble in ammonium hydroxide, and 
in sodium thiosulphate, "hypo." It is much more sensi- 
tive to light than the silver nitrate, and hence for photo- 
graphic work the latter salt is generally converted into the 
chloride, or bromide, which is even more sensitive. It is 
believed that the light gradually converts this compound 
back into metallic silver, which is insoluble in the " hypo," 
while the unchanged portions of silver chloride are dis- 
solved out and the paper thus de-sensitized. 

Experiment 144. — To about 1 cc. of a solution of silver nitrate 
add a few drops of hydrochloric acid. Notice the appearance of the 
precipitate that forms. What is it? Write the reaction. To a por- 
tion of it add a little ammonium hydroxide and shake it. What 
results ? To another portion add a solution of u hypo " and state the 
results. 

28. Silver Chromate, Ag 2 Cr0 4 . — This is a blood-red 
powder obtained as a precipitate when potassium chro- 
mate is added to a solution of silver nitrate. 

Experiment 145. — Prepare the chromate as indicated, and note 
its appearance. 

29. The formation of silver chloride and the chromate, 
with their characteristic appearance and the ready solu- 
bility of the former in ammonia, serve to distinguish a 
solution of silver, and may be used as tests. 



THE COPPER-SILVER GROUP 243 

Exercise. — Write the reactions, showing the preparation of sil- 
ver nitrate, silver chloride, and the chromate ; also silver bromide and 
iodide, from silver nitrate with potassium bromide, and with sodium 
iodide. 

30. Photography. — At the present time almost all young 
people take more or less interest in this wonderful art. 
The first experiments along this line were made as early 
as 1727, but they were nothing more than what the student 
has done in the first part of Experiment 143, and the print 
soon disappeared. From that time to this many different 
plans have been tried, but we can only notice briefly that 
used at present. 

31. Preparation of the Plates. — The plates upon which 
the negatives are made are prepared as follows : for the 
most sensitive plates, potassium or ammonium bromide 
with gelatin and silver nitrate added is dissolved in water 
and heated to boiling. Thus the silver is converted into 
silver bromide : — 

KBr + AgN0 8 = AgBr + KN0 3 . 

An excess of water is added, and the potassium nitrate 
formed is readily washed away. This gelatin emulsion, 
as it is known, is poured upon glass plates and allowed to 
harden ; they are then ready for use. 

32. Exposure and Developing. — As previously stated, 
when such plates are exposed to light, the silver salts are 
decomposed. In the camera the exposure is so brief that 
the decomposition is only partial; when, however, the 
plate is put into the developer, this solution continues 
the action begun by the light. Hence those portions of 
the plate which have received the most light have the 
larger amount of the silver salts decomposed, and are dark 
in color. If allowed to remain in the developer long 



244 MODERN CHEMISTRY 

enough, all the silver would be reduced, and the plate 
would be uniformly dark. 

33. Fixing. — When it is seen by examination that the 
development has proceeded long enough, the plate is rinsed 
in water and placed in the fixing bath. This is a solution 
containing sodium thiosulphate, which is an excellent 
solvent for many silver compounds. The fixing bath soon 
removes from the gelatin film the silver bromide or chlo- 
ride that remains unaffected by the light or by the devel- 
oper. The plate is thus cleared or fixed, and is no longer 
sensitive to light. As the lights and shadows are all 
reversed, it is called a negative. After thorough washing 
it is allowed to dry, when it is ready to be used in making 
prints. 

34. Printing. — Various kinds of paper are now used 
for making prints, among them being the solio, velox, 
platinotype, carbon, and blue print. The first and last of 
these require the least skill. Solio has a sensitized film 
of silver chloride ; in printing, this is placed against the 
film side of the negative, which causes the objects to 
appear in the picture in their natural position. As the 
dark portions of the negative transmit the fewer light 
rays, the picture appears as a positive, or like the original 
as to high lights and shadows. The advantage of solio 
is in the fact that it is only moderately sensitive, and 
that it readily shows when it has been exposed long 
enough. More sensitive papers, such as the velox, are 
like the gelatin plates in that they show no image until 
treated with a developer. Solio prints require toning, 
and all varieties need fixing by some method or other. 
In the platinotype papers, a compound of platinum is 
used which yields the d.ark appearance now so much 
admired. 



THE COPPER-SILVER GROUP 245 

35. Solio papers cost so little that it would be easy for 
a class to make some experiments along this line. Let 
any of the pupils who may have them bring in some of 
their negatives and printing frames, and do some work of 
this kind. 

36. Blue prints* are the simplest of all, are cheap, and 
yet for landscapes often give most excellent effects. 
They possess the advantage of requiring no toning or 
fixing except such as is secured by thorough washing. 
Place the paper under the negative in direct sunlight, 
and allow it to remain until the high lights begin to look 
somewhat muddy in appearance ; then put into a basin of 
water with the printed side down. Allow the print to 
remain there until the light portions are quite clear, then 
wash for ten minutes in running water. The beauty of 
these prints will be enhanced by leaving a pure white 
border around the picture ; this may be secured by using 
a black mat between the negative and print so as to cover 
the portion which it is desired to have white. 



* If he desires, the instructor may prepare his own blue print paper. 
Make 

Solution A 

Oxalate of iron and ammonia . . 1 g. 
Water 10 ec. 

Solution B 

Potassium ferricyanide . . . 1 g. 

Water 10 cc. 

When ready for use mix A and B in a dark room, and apply to the paper 
with a brush ; or, the paper may be floated in the solution. This must 
be used within a day or two after it is prepared, as it does not keep 
well. A few drops of a 10 per cent solution of potassium bromide 
added to A and B above will render the keeping qualities of the paper 
much better. 



246 MODERN CHEMISTRY 

Gold: Au = 197 

37. Occurrence. — From the fact that gold occurs free, 
it has been known from the earliest antiquity. It is 
widely distributed over various portions of the earth and 
usually occurs in fine grains and nuggets disseminated 
through the rocks. These are gradually disintegrated 
and brought down by rains and streams in the form of sand 
and gravel, with which the gold is associated. The best- 
producing gold regions are those of the western part of 
the United States, Australia, Southern Africa, and the 
Klondike. Gold also occurs in quartz veins deeply buried 
in the earth's strata. 

38. Methods of Mining. — The original method consisted 
simply in cradling or panning the sand and gravel ; thus 
the nuggets and larger grains find their way to the bot- 
tom, while the lighter stone and earthy matter is washed 
out. By this method only the larger particles are saved. 
Placer mining consists in washing the gold-bearing sand 
down through sluices, along the bottom of which are 
arranged pockets of mercury, or over plates of copper 
amalgamated with mercury. This readily amalgamates 
with the gold, and the other portions are carried away 
by the current. The gold amalgam thus obtained is 
heated in retorts, by which the mercury is vaporized, 
leaving the gold behind. The vapors of mercury are 
conducted into cold chambers where they are condensed, 
so that very little loss occurs. Hydraulic mining differs 
from the above only in that streams of water are directed 
with great force against the loose rock and cliffs over- 
hanging, washing them down into the sluice-ways. 

39. Vein Mining. — Vein mining differs from placer 
mining in that the latter is surface mining, while in the 



THE COPPER-SILVER GROUP 247 

former the ore is taken from greater or less depths, usually 
from quartz veins ; hence it is sometimes called quartz 
mining. Gold sometimes occurs in combination with 
iron in pyrites, and it is then obtained by the wet or 
ehlorination process. The ore is roasted, then moistened 
and treated with chlorine, obtained usually from bleach- 
ing powder. The chlorine dissolves the gold, forming gold 
chloride, AuCl 3 . This is now dissolved out and ferrous 
sulphate added, which precipitates gold in the metallic 
condition, as seen in the following reaction : — 

6 FeS0 4 + 2 AuCl 3 = 2 Au + 2 Fe 2 (S0 4 ) 3 + Fe 2 Cl 6 . 

40. Cyanide Process. — Potassium cyanide is an excel- 
lent solvent for gold, and at the present time is used 
extensively in separating it from its ores. The process 
is valuable where the gold occurs in a finely divided form; 
another advantage is that the ore does not need the roast- 
ing that is necessary in the ehlorination process. After 
the gold-bearing quartz has been finely crushed, it is 
treated with a solution of potassium cyanide in water. 
The gold is dissolved out, thus : — 

4 Au + 8 KCy + 2 + 2 H 2 = 4 KAuCy 2 + 4 KOH. 

41. The oxygen shown in the reaction is derived from 
the air, and it has been found that, unless the surface of 
the ore is left well exposed, the process is not satisfactory. 
The double cyanide of gold and potassium thus obtained 
is treated with zinc, which precipitates the gold, as shown 
in the reaction : — 

2 KAuCy 2 + Zn = K 2 ZnCy 4 + 2 Au. 

42. There is always some zinc left in a more or less 
finely divided form which cannot be separated mechani- 
cally from the gold ; hence, when melted down the metal 



248 MODERN CHEMISTRY 

is seldom over 80 per cent pure. For this reason some 
companies prefer to deposit the gold by electrolysis upon 
lead terminals. By this method, after oxidizing the lead 
in cupels, the gold remains in a very pure form. 

43. Characteristics. — Gold is a bright yellow metal, 
which, seen in light reflected several times, looks red. It 
is so soft that for ordinary purposes it must be alloyed 
with some other metal ; it is heavy, is not affected by the 
oxygen of the air at any temperature, is very ductile and 
malleable. Advantage is taken of this property in ham- 
mering out the metal into gold leaf, the thickness of which 
is not over 3-00W P ar ^ °^ an i n °h> 1500 of which sheets 
are necessary to make one as thick as ordinary note paper. 
Pure gold is not affected by single acids, but is readily 
attacked by aqua regia, forming gold chloride, AuCl 3 . 
However, if richly alloyed with several other metals, it 
becomes soluble in single acids. 

44. Uses. — These are too well known to need specifi- 
cation. In the arts gold leaf has numerous uses, such as 
in making display signs, covering high grade moldings, 
for filling teeth, etc. 

SUMMARY OF CHAPTER— COMPARATIVE STUDY 

Copper, Silver, Gold. 

Histories — Wherein are they similar ? — Why ? 
Occurrence — In what forms ? 
Most productive regions. 
Some important ores. 
Various forms of gold mining — Description. 
Reduction of the ores. 

Special plans for copper. 

Special processes for gold reduction. 

Chlorination and cyanide. 
Special plans for separation of silver. 
Pattison's and Parke's. 



THE COPPER-SILVER GROUP 249 

Comparison of the three metals as to 

a. Color. 

b. Density. 

c. Melting point. 

d. Permanency in the air. 

e. Malleability. 
/. Conductivity. 

g. Solubility in acids. 
Uses of the metals. 

a. Important alloys. 

b. Other uses — Why so used. 
Compounds — Most important. 

Of Copper — The Sulphate — Commercial name and formula. 

How obtained. 

Characteristics and uses. 
Of Silver — The Nitrate — Commercial name and formula. 

How prepared. 

Appearance and uses — - Why so used ? 

AVhat other compounds prepared from this one? How? 
Special points. 

Meaning of the terms blister copper, matte, concentration, 

calcination, converting. 
Describe method of determining the composition of brass. 
Meaning of terms cupel, cupellation. 

Describe experiment illustrating principles of photography. 
Method of sensitizing photographic plates. 
Chemistry of the developing and fixing of negatives. 
Reactions showing the preparation of CuS0 4 , CuO, AgX0 3 , 
AgCl, AgBr, AgL 



CHAPTER XXI 

ZINC, CADMIUM, MERCURY 
Zinc : Zn = 65 

1. History. — Brass, an alloy of copper and zinc, has 
been known for centuries, but it was formerly made by 
fusing together copper and a mineral called calamine, 
which we now know is an ore of zinc. It was not until 
about the close of the seventeenth century that zinc was 
recognized as a distinct metal and its characteristics care- 
fully determined. 

2. Ores of Zinc. — Zinc occurs abundantly in many parts 
of the United States and Europe. In Missouri the mines 
of Joplin and Webb City are the best known. There 
thousands of tons are produced annually. Kansas also 
yields a considerable quantity. The ore most generally 
found in these states is the sulphide, ZnS, known as zinc 
blende. By the miners it is called " jack," or in its purer 
forms " rosin jack," because of the general resemblance of 
a broken specimen of the ore to rosin. In New Jersey the 
ore franklinite is the most abundant. It is a mixture of 
zinc oxide and ferric oxide, ZnFeO, Fe 2 3 . Other sections 
yield the carbonate, ZnC0 3 , known as smithsonite. It is 
said that the metal is sometimes found pure in Australia. 

3. Reduction of the Ores. — The general method em- 
ployed in the reduction of the greater number of metallic 
ores is used in the case of zinc. Thej^ are first ground 
fine and roasted. This not only drives out certain volatile 

250 



ZINC, CADMIUM, MERCUBY 251 

impurities, such as arsenic, but converts the ore into the 
oxide, ZnO, the most convenient form for the next step. 
The reaction that takes place when the ore is roasted may 
be seen from the following : — 

ZnS + 30 = ZnO + S0 2 , 

ZnC0 3 + heat = ZnO + C0 2 . 

4. The oxide thus obtained is mixed with powdered 
coke and heated red hot in earthen cylinders about 4^ 
feet long, placed horizontally over one another. The zinc 
is thus reduced to the metallic form, and at the tempera- 
ture obtained is vaporized. The vapors pass out into 
conical-shaped earthen condensers attached to the outer 
end of the retorts, where they liquefy. Twice in twenty- 
four hours these condensers are " tapped " and the molten 
zinc drawn off and run into molds. The chemical change 
taking place is a familiar one : — 

ZnO + C = Zn + CO. 

The retort is shown by R in the figure and the con- 
denser by 0. The condensers are readily detached, and 
1 when the retorts have been charged or filled with the 
mixed ore and coke they are 
again attached and luted on ,^^ \ R 

nearly air-tight with clay. ^ 
When in operation there is 

usually some escape of vaporized zinc with other gases, 
and these, in burning at the mouth of the condensers, 
give a beautiful display of colors, yellow and blue and 
white, which, especially at night, is exceedingly striking. 

5. A single charge requires about twenty-four hours 
for complete reduction, and as the workmen are usually 
paid by the amount of metal they " draw off " they gen- 




252 MODERN CHEMISTRY 

erally work twenty-four hours successively, and then are 
off during the next twenty-four. The zinc obtained in 
this way is more or less impure ; it almost always contains 
some cadmium, and usually some arsenic, and is known 
as " spelter." 

6. Characteristics of Zinc. — Zinc is a bluish white 
metal of moderately low melting point, about 420° C; it 
tarnishes but slightly in the air, and then only upon the 
surface. At a temperature slightly above the melting 
point it burns with a brilliant, bluish white flame, and if 
a jet of oxygen be directed upon it the light is almost 
dazzling. 

Experiment 146. — Examine a piece of zinc and note its color, 
malleability, hardness, and tendency to oxidize. Test also its melting 
point by heating a small piece on charcoal with the blowpipe. Try 
it also with the oxidizing flame and note the deposit upon the char- 
coal, both when hot and when cold. State the results. 

Experiment 147. — To learn the solvents for zinc. Try a small 
piece of the metal in a test-tube with hydrochloric acid. How is it 
affected? What gas is obtained? What proof can you offer? Write 
the reaction. 

In the same way try nitric acid, and compare results with the 
above. * 

Into each of two test-tubes put a small piece of zinc. To one add 
about a cubic centimeter of copper sulphate solution, and cover the 
other with water. After a few moments, to each add a little sulphuric 
acid. Is there any difference in the rapidity of the chemical action in 
the two cases ? Why ? 

Experiment 148. — Sift some zinc dust through a wire sieve of 
fine mesh upon a Bun sen burner flame and note the results. 

7. Further Characteristics of Zinc. — Ordinary com- 
mercial zinc as it comes from the smelter is brittle, but if 
it is heated to something over 120°, and then rolled into 
sheets or drawn into wires, it is found to be malleable, and 
will remain so. As it approaches the melting point, how- 



ZINC, CADMIUM, MERCURY 253 

ever, it again becomes brittle, and may be ground into a 
powder known as zinc dust. It is of medium density, 
being a little lighter than iron, is not magnetic, and when 
chemically pure is but slightly soluble in dilute acids. 
When impure, or if in contact with some other metal, as 
copper or platinum, the solution is rapid. 

8. Uses for Zinc. — In the metallic form zinc is used 
extensively in many varieties of galvanic batteries, also as 
linings for refrigerators, bathtubs, and for various other 
domestic purposes. One of its most important applica- 
tions is in coating or " galvanizing " iron wire and other 
forms of iron as a protection from moisture. Galvanized 
iron is prepared by thoroughly cleansing the iron to be 
coated, heating it, and plunging it into a bath of molten 
zinc until a thin covering of the latter metal adheres. 
There are also three important alloys : — 

Brass : consisting of zinc and copper in varying propor- 
tions ; 

Bronze : zinc, copper, and. tin ; 

German silver : zinc, copper, and nickel. 
In the chemical laboratory zinc is frequently used : in 
making hydrogen; in reducing ferric compounds to the fer- 
rous condition ; and for precipitating various metals from 

their solutions. 

Compounds of Zinc 

9. Zinc Sulphate, ZnS0 4 , 7 H 2 0. —White Vitriol. —This 
is a white crystalline salt which has been obtained in the 
preparation of hydrogen by treating zinc with sulphuric 
acid. It is very soluble in water, and is used mainly for 
calico printing. It has a bitter, astringent taste. 

10. Zinc Chloride, ZnCl 2 . — This is a white solid, ob- 
tained when zinc is dissolved in hydrochloric acid. It has 
great affinity for water, and is, therefore, often used in 



254 MODERN CHEMISTRY 

chemistry as a drying agent. It is also frequently used as 
a soldering solution, but as it is poisonous, serious results 
have sometimes followed its use in soldering tin cans con- 
taining fruits and other food products. 

11. Zinc Hydroxide, Zn(OH) 2 . — This compound of zinc 
may be studied in the following experiment : — 

Experiment 149. — To a few cubic centimeters of a solution of 
any zinc salt, as the chloride or sulphate, add a few drops of ammo- 
nium hydroxide. What are the results ? Add more ammonia ; does 
the precipitate dissolve? Describe the precipitate, Zn(OH) 2 , that 
formed, and write reaction. In the same way prepare a little zinc 
hydroxide by using a solution of caustic soda or potash instead of am- 
monia as above. Test a portion of the precipitate with hydrochloric 
acid ; does it dissolve ? Write the reaction. 

12. Zinc Sulphide, ZnS. — Many characteristics of zinc 
sulphide may be discovered from the following experi- 
ment : 

Experiment 150. — To a few cubic centimeters of a solution of 
some zinc salt add two or three drops of ammonium sulphide. De- 
scribe the precipitate that forms. It is zinc sulphide. Test its solu- 
bility in dilute hydrochloric or nitric acid. 

13. Zinc Oxide, ZnO. Zinc White. — This was the white 
deposit formed on charcoal when the zinc was heated by 
the oxidizing flame. It is now used extensively as a sub- 
stitute for white lead in painting, and is preferable in 
localities where much coal is used as fuel, because of the 
discoloration of lead compounds by the considerable quan- 
tities of hydrogen sulphide found in coal smoke. 

* Cadmium, Cd = 112 

14. Supply. — Cadmium is a rare element, discovered 
about 1817. It received its name from a Greek word, 

* This is an unimportant element, and its study may be omitted, if 
desired. 



ZINC, CADMIUM, MERCURY 255 

kadmeia, an ore of zinc, now known as calamine, with 
which cadmium is usually associated. Our present supply 
is obtained mostly from zinc ores, with which it is found, 
in the form of a sulphide, CdS, called greenockite. 

15. Reduction of the Ore. — In smelting cadmium-bear- 
ing zinc ores, they are first roasted in retorts, where both 
sulphides are converted into oxides, thus : — 

CdSj _(CdO 

These oxides are then mixed with coke or charcoal and 
again heated, when the usual reduction takes place : — 

""l + io.ff+.co. 

ZnOj [Zn 

To separate the two metals thus obtained they are dis- 
solved in hydrochloric acid, and the solution treated with 
rods of zinc, by which the cadmium is reduced to the 
metallic form, thus : — 

and 

16. Appearance and Characteristics. — Cadmium is usu- 
ally marketed in the form of small rods, 8 or 10 inches 
in length. It is a white metal, closely resembling tin, and 
is of about the same hardness, but it has a melting point 
not very different from lead, 315°, and boils at 860°. 
Cadmium tarnishes slowly in the air, becoming coated 
with a very thin covering of yellow oxide. It is malleable 
and ductile, and when bent, like tin, gives a creaking sound. 
With mercury it forms a silvery white amalgam which 



256 MODERN CHEMISTRY 

soon becomes hard and brittle. It is easily soluble in 
nitric acid, less so in hydrochloric and sulphuric acids. It 
is but little used in the arts, though it has been applied 
somewhat as a filling for teeth ; but as a cadmium amal- 
gam gradually turns dark, it has not found favor with 
dentists. 

Compounds of Cadmium 

17. Cadmium Nitrate, Cd(N0 3 ) 2 . — This is a white salt 
obtained when the metal is dissolved in nitric acid. 

18. Cadmium Sulphide, CdS. — This is a yellow powder 
used in oil and water colors. Artificially, it is obtained 
when a current of hydrogen sulphide is passed through a 
solution of any cadmium salt. It resembles the sulphides 
of arsenic and tin, As 2 S 3 and SnS 2 . Unlike the arsenic, 
however, cadmium sulphide is not soluble in ammonium 
carbonate, and unlike the tin, is insoluble in yellow ammo- 
nium sulphide. 

Exercise. — Write reactions showing the formation of cadmium 
nitrate, chloride, sulphate, and sulphide. 

Mercury : Hg = 200 

19. Historical Facts. — Mercury was one of the seven 
elements known to ancient chemists, and by them was 
dedicated to the god from which it received its name. Its 
symbol is taken from the Greek word, hydrargyrum, by 
which name it was also known. This term means water 
(or liquid^ silver. Similarly, at the present time it is 
spoken of as quicksilver. By Geber, the famous alchemist 
of the eighth century, mercury and sulphur were regarded 
as the two elements from which all metals could be made. 
He claimed that any one knowing the proper proportions 
could prepare any of the noble metals from these two. 



ZIXC, CADMIUM, MERCURY 



20. The Source of Supply. — The commercial supply of 
mercury comes from its chief ore. cinnabar, or vermilion, 
HgS. This is an exceedingly heavy, brick-red mineral, 
found in Spain, India, Bavaria. California. Mexico, etc. 

Experiment 151. — Xear one end of a piece of hard glass tubing 
place a little vermilion, HgS. as much as will remain on the point of a 
knife-blade. Xow, with this end down, hold in a slanting position in 
the Bmisen burner and heat strongly. Xotice the formation on the 
upper, cooler portion of the tube. What gas. detected by its odor, is 
given off from the upper end of the tube? Xame the two products 
resulting from the heating of mercuric sulphide. Compare with the 
preparation of oxygen from mercuric oxide. 

21. Reduction of the Ore. — This experiment illustrates 
the reduction of cinnabar in the preparation of mercury 
for commerce. The ore is placed 
upon shelves in an oven over a 
furnace (see Fig. ob^). Hot blasts 
of air flow up through the shelves. 
oxidizing the sulphur to sulphur 
dioxide, and at the same time 
vaporizing the mercury. These 
gases pass out together into cool chambers, where the mer- 
cury condenses, while the sulphur dioxide escapes. As 
thus obtained the mercury is more or less impure. It is 
purified first by being strained through porous leather or 
chamois skin, and then distilled at moderate temperatures. 

22. Characteristics. — Mercury is the only metal that is 
liquid at ordinary temperatures. At 39° below zero it 
becomes a solid, and in that condition possesses some 
of the properties of lead. It has about the same color, is 
malleable, and soft enough to be cut easily. Mercury is 
a silver-white metal, which does not tarnish in the air. but 
which slowly vaporizes at all temperatures. 




Fig. 55. 



258 MODERN CHEMISTRY 

23. Amalgams. — The most remarkable property of 
mercury is its power of dissolving many of the metals 
and forming with them what are known as amalgams. 
If the mercury be largely in excess, the other metal dis- 
appears as a lump of sugar does in a cup of tea ; if a 
smaller proportion be used, the mercury simply combines 
with the outer portions of the other metal, changing more 
or less its appearance and general properties. There are 
two methods of forming amalgams. 

24. a. By bringing metallic mercury into contact with 
a metal perfectly clean. If this is broken up into small 
pieces, or in the form of dust or filings, and is then heated 
with the mercury, the amalgamation takes place quickly. 

Experiment 152. — Into a few drops of mercury in an evaporating 
dish, put a perfectly clean strip of zinc. After a few moments, ex- 
amine it; has it changed in appearance? Bend it; has it changed in 
properties ? Try in the same way a five-cent piece, a penny, a nail, or 
any other convenient metals. Be careful, however, of any gold rings, 
as mercury amalgamates very readily with gold. 

25. h. The second general method is by immersing the 
metal to be amalgamated in a solution of some salt of 
mercury. Try in this way the following : — 

Experiment 153. — Put into a beaker, or evaporating dish, a few 
cubic centimeters of a solution of mercurous nitrate, Hg 2 (N0 3 ) 2 . 
Immerse in it a brass pin, or a thimble, a copper penny, a key ring, 
etc. After remaining a few minutes, they may be removed and 
rubbed a little, if dull in appearance. State which have been 
amalgamated. 

26. This second method is employed frequently by 
street fakirs as a means of " silver plating." They pre- 
pare the solution by dissolving mercury in nitric acid 
and then adding some coloring matter. This very rapidly 



ZINC, CADMIUM, MEBCUBY 259 

" plates " certain metals, but the amalgamated articles 
retain their brilliancy but a short time. 

27. Solvents for Mercury. — The best solvent for mer- 
cury is nitric acid, which attacks the metal even at or- 
dinary temperatures. When heated, sulphuric acid also 
dissolves it, with the formation of sulphur dioxide gas. 
Compare this with the preparation of sulphur dioxide as 
given on page 177, section 15. 

28. Uses for Mercury. — Mercury is employed exten- 
sively in the manufacture of thermometers and barom- 
eters ; in the laboratory it is often used instead of 
water in the pneumatic trough for collecting such gases 
as are soluble in water, especially when they are desired 
perfectly free of air. Large quantities are also used in 
placer mining of gold and silver (see page 246). In the 
form of amalgams it is used with various other metals for 
filling teeth ; with tin or silver for the backs of mirrors, 
for rendering zinc plates to be used in batteries less solu- 
ble in acids, and sometimes for amalgamating surfaces 
which are to be silver plated. This is done because silver 
seems to adhere better to a surface which has been thus 

treated. 

Compounds of Mercury 

29. Like several other metals, mercury forms two series 
of compounds, the mercurous and mercuric. 

30. The Nitrates, Mercurous, Hg 2 (N0 3 ) 2 ; Mercuric, 
Hg(N0 3 ) 2 . — These may be prepared by treating mercury 
with nitric acid ; for the former, using dilute acid with 
the mercury in excess ; for the latter, concentrated, with 
the acid in excess. Mercuric nitrate is a white salt of fine 
silky crystals, soluble in water. In dissolving it yields at 
the same time a yellowish powder, known as basic nitrate, 
having the formula HgN0 3 , Hg(OH) 2 . Mercurous ni- 



260 MODERN CHEMISTRY 

trate is of a pale yellow color, almost white. It usually 
occurs in crystals larger than those of the mercuric nitrate 
and is soluble in water. Both are used in the laboratory, 
and occasionally for the preparation of other compounds 
of mercury. 

31. The Chlorides, Mercurous, Hg 2 Cl 2 ; Mercuric, HgCl. r 
— The former is known as calomel, the latter as corrosive 
sublimate, Mercurous chloride may be prepared by add- 
ing to mercurous nitrate, hydrochloric acid, whereupon it 
falls as a heavy white precipitate. On a large scale it is 
manufactured by thoroughly mixing in the proper propor- 
tions mercuric chloride and mercury, heating them strongly 
to vaporize, wdiereupon they combine and are condensed in 
cold chambers. Calomel is a white, flour-like substance, 
insoluble in water. It is used largely in medicine. 

32. Mercuric chloride is prepared by subliming, as de- 
scribed above in making calomel, a mixture of mercuric 
sulphate and common salt. It is a white, crystalline salt, 
somewhat soluble in water, and very poisonous. It is used 
in the laboratory as a reagent, is a constituent of some 
vermin exterminators, and has frequent use in surgery as 
an antiseptic. 

33. Mercuric Oxide, HgO. — This orange-red salt, com- 
monly known as red precipitate, is prepared by heating 
mercuric nitrate for a considerable length of time. It is 
used sometimes for preparing small quantities of oxygen, 
and in some quantitative determinations in the laboratory. 

34. Mercuric Sulphide, HgS. — As an ore it is known as 
cinnabar, but the artificial product is sold under the name 
vermilion* It is of a bright scarlet color, and is used 
in making tube paints and in coloring sealing-wax. As 
ordinarily prepared in the laboratory, it is black, but under 
certain conditions is obtained in varying shades of red. 



ZINC, CADMIUM, MERCURY 



261 



Experiment 154. — To prepare certain compounds of mercury. Put 
a drop of mercury into a test-tube and add about a cubic centimeter of 
dilute nitric acid, warm gently, and after a few minutes, or when the 
action has ceased, decant the solution and boil it nearly dry in an 
evaporating dish. Now add a few cubic centimeters of water and pour 
into three test-tubes. To the first add a little potassium iodide ; to the 
second, hydrochloric acid ; to the third, ammonia. Notice the color of 
the precipitate in each case. Write the reactions in the first two, and 
state what compound is formed. Tabulate results as follows : — 



Hg 2 (XO.,) 2 



Hg(NO s ) 2 



KI 






HC1 






NH 4 OH 







You should have prepared mercurous nitrate by the above treat- 
ment of mercury with nitric acid. 

To another drop of mercury in a test-tube add some strong nitric 
acid, and warm until the mercury is all dissolved. Transfer to an 
evaporating dish and boil nearly dry, then add a few centimeters of 
water. You should now have a solution of mercuric nitrate. Divide 
into three parts and treat with the same three reagents that you used 
with the mercurous nitrate, and tabulate the results. 

35. From the above experiments, it will be seen that 
the two series of mercury salts may be easily distinguished 
by the precipitates which they form with different reagents. 

Experiment 155. — Let the student be given some mercurous and 
mercuric solutions, and have him determine what each is. 

Exercise. — Write out the reactions that take place in preparing 
the two nitrates, mercurous chloride, mercurous and mercuric iodide, 



262 MODERN CHEMISTRY 

and mercuric sulphate. Before attempting the last, unless you know 
the results, put into a test-tube a small drop of mercury, add a little 
strong sulphuric acid, and heat until some familiar gas is produced. 

COMPARATIVE STUDY 

Zinc and Mercury — Early history. 

Ores of these metals — Most important of each — Localities where 

found. 
Plan of reduction. 

Wherein are they alike? 

How different ? 

Why is carbon not necessary for the mercury ? 

Reactions for each. 

Description of furnaces. 
Comparison of the two metals in 

a. Color. 

b. Melting point. 

c. Density. 

d. Ease of oxidation. 

e. Malleability. 
/. Conductivity. 
g. Solubility. 

Special properties of each. 

Brittleness of zinc at certain temperatures. 
Condition of mercury at low temperatures. 
Power of forming amalgams. 

Names of metals which will amalgamate and of those 
which will not. 
Two methods of making amalgams. 
Important uses of each metal. 
Alloys of zinc. 
Amalgams of mercury. 
Other uses. 
Compounds. 

The oxides — Appearance and use of each. 
The chlorides — One of zinc, two of mercury. 

Preparation of each — Use — Commercial name. 
Two classes of mercury compounds — Methods of distin- 
guishing them. 



CHAPTER XXII 

ALUMINUM AND ITS COMPOUNDS 

Aluminum : A\ = 27 

1. Abundance. — This metal was first isolated about 
1827, being reduced by metallic sodium. For some 
years all that was used in the arts was prepared' by strongly 
heating aluminum chloride, and passing the vapors into 
which it was converted over sodium. The reaction may 
be represented thus : — 

A1C1 8 + 3 Na = Al + 3 NaCl. 

By this method about three pounds of sodium were re- 
quired for the preparation of a single pound of aluminum, 
and the cost was about one dollar an ounce. 

2. No metal occurs more abundantly than aluminum, 
and but one or two non-metallic elements are more widely 
distributed. It forms a large per cent of feldspar and of 
various other rocks, and consequently, from their decom- 
position, of all clays. 

3. The Commercial Supply. — It is evident that all that 
is needed to insure a large output of aluminum is a cheap 
process of reducing it from its natural compounds. Various 
methods have been patented, but none has, as yet, brought 
aluminum within the reach of all, although its market 
value now is only about $1.50 per pound. Perhaps the 
most satisfactory plan yet adopted is the following : a is a 
large crucible lined with some infusible substance, like 

263 



264 



MODERN CHEMISTRY 




Fig. 56. 



graphite ; c is a bundle of carbon rods. The crucible 
and carbons are made the kathode and anode from the 

dynamo d. Into the crucible 
is put cryolite, a compound of 
aluminum and sodium fluor- 
ide, Al 2 Na 6 F 12 , mixed with 
bauxite, or aluminum oxide, 
A1 2 3 . The former compound 
has a very low melting point 
and serves as a flux. When 
the mixture of the two min- 
erals has been fused, a powerful current is passed through, 
and the bauxite alone is decomposed. We may represent 
this by the simple equation 

A1 2 3 = Al 2 + 3 O. 

The metallic aluminum collects at the bottom of the 
crucible and may be drawn off. 

4. Characteristics. — Aluminum is a silvery white metal, 
having a density of only 2.6, or about one-third that of 
iron. It is very tenacious, ductile, malleable, and sonorous. 
Its melting point is only moderately high, 700°. It is 
permanent in the air and a good conductor of electricity. 
Aluminum is strongly electropositive in its character, and 
may be used to reduce various other metals from their 
compounds, just as zinc will reduce lead, tin, and others. 

5. Uses. — As the output of aluminum has increased* 
and the price cheapened, its uses have rapidly multi- 
plied. 

6. The metal will take a very high polish, and as it is 
permanent in air it is now used extensively in the manu- 
facture of ornaments and novelties, for which, in the past, 
silver alone had been employed. 



ALUMINUM AND ITS COMPOUNDS 265 

7. A more valuable use of aluminum is in the place of 
copper in electric circuits. Because of its lightness, an 
aluminum wire, much larger in cross-section than any- 
copper wire used for this purpose, may be employed with- 
out increasing the weight in a given length. This fact 
will nearly suffice to offset the lower coefficient of conduc- 
tivity of the newer metal, and makes its carrying capacity 
not very different from that of copper. Another practical 
use of aluminum is in the manufacture of cooking utensils. 
It is claimed that these vessels will prevent the scorching 
of liquid foods, such as milk. The metal is also used in 
many places where iron has heretofore been employed, 
having the advantage of great tenacity. In time, if the 
cost of production is sufficiently decreased, it may find 
extensive use in shipbuilding. In the form of alloys, 
with varying proportions of copper, it is much used. 

Compounds of Aluminum — Native 

8. Kaolin. — This is one of the most valuable natural 
compounds of aluminum ; it is the silicate, Al 4 (Si0 4 ) 3 , 
almost pure. Ordinary clays contain, in addition to this, 
iron compounds and other foreign materials. Kaolin, 
when heated with a kind of rock called feldspar, melts 
and forms a semi-translucent mass, used in making various 
kinds of porcelain wares. 

9. Emery. — In the form of a rock or mineral, emery is 
known as corundum. It is the oxide of aluminum, A1 2 3 , 
and is extremely hard. It is used in the form of emery 
wheels, and in other ways, for polishing and for grinding 
and sharpening tools. 

10. Precious Stones. — The Oriental ruby, the emerald, 
sapphire, and many other stones prized for their beauty 
are merely aluminum oxide, A1 2 3 , crystallized with some 



266 MODERN CHEMISTRY 

silica, Si0 2 , and differ in almost no other respect from 
emery or corundum, which is uncrystallized. These jewels 
contain in addition very small quantities of some foreign 
substance, such as iron, chromium, or copper, which im- 
parts the colors for which they are valued. The true 
or Oriental ruby and other gems differ from the spinel 
or ordinary ruby and emerald in that the latter contain in 
addition to the aluminum oxide, certain other compounds, 
which render the stones much less valuable. 

Compounds of Aluminum — Artificial 

11. Alums. — The term alum is applied to a large num- 
ber of salts, known as double sulphates. They all contain 
two metals, or an equivalent. Thus, common alum is 
potassium aluminum sulphate, K 2 A1 2 (S0 4 ) 4 24 H 2 0. By 
inspection, it will be seen that this is simply potassium 
sulphate, K 2 S0 4 , crystallized with aluminum sulphate, 
A1 2 (S0 4 ) 3 . They all crystallize in octahedrons, some- 
times singly, more usually piled in masses one upon 
another. 

Experiment 156. — Put about 50 cc. of water into a beaker, warm 
it somewhat, and add powdered alum as long as any will dissolve. 
Now, pour the saturated solution into a good-sized test-tube, and sus- 
pend therein a string with a small knot at the end. Allow it to stand 
several hours and note the shape of the crystals that form. By adding 
a strong chromium solution to the alum, delicately colored violet crys- 
tals may be obtained. 

12. Common alum is prepared by burning a shaly rock 
which contains a compound of aluminum, moistening it, 
and exposing it to the air. The aluminum compound is 
thus converted into a sulphate. A potassium salt is then 
added to the solution, whereupon the alum crystallizes out. 
If the ammonia water from gas factories be used instead 



ALUMINUM AND ITS COMPOUNDS 267 

of the potash salts, ammonia alum is obtained, represented 
by the formula (NH 4 ) 2 A1 2 (S0 4 ) 4 24 H 2 0. This is used 
very largely instead of the potash alum. 

13. The term alum is also applied to many double 
sulphates which contain no aluminum, as, for example, 
K 2 Cr 2 (S0 4 ) 4 , potassio-chromic alum. In such cases the 
compound is designated by the names of both of the metals 
entering into the compound. By studying the formulae 
for the alums mentioned above, it will be seen that the 
first metal is always univalent and the second usually triv- 
alent. If, then, we represent the univalent metals by 
M and the trivalent by i£, we may write as the general 
formula for the alums, M 2 R 2 (S0 4 ) 4 , in which M is usually 
potassium, sodium, or ammonium, and R aluminum, iron, 
or chromium. 

Exercise. — Write formulae for sodium alum, ammonio-ferric alum, 
potassium alum, sodio-chromic alum, potassio-f erric alum, and ammonio- 
chromic alum. 

14. Alum possesses a sweetish, astringent taste, and 
upon heating readily gives up its water of crystallization. 
In doing so it crumbles to an opaque mass, and in this 
form is known as burnt alum. 

15. Uses. — In medicine, alum is used as an astringent. 
It checks bleeding by contracting the tissues, and in the 
form of burnt alum it serves as a mild caustic agent, 
especially for ulcerations of the mouth. In the arts alum 
is used largely as a mordant, that is, to fix the color in 
dyeing cloth. 

Experiment 157. — In a solution of logwood heat two pieces of 
cloth for several minutes, or until both are well colored. Now, re- 
move and allow them to dry. Then immerse one in a strong solution 
of alum and let it stand a few minutes ; remove, and when dry wash 
both in water. Which retains its color the better ? 



20)8 MODERN CHEMISTRY 

16. Alum is also used frequently as an adulterant for 
baking powders, sometimes in very considerable quantities. 

Experiment 158. — Examine a number of specimens of baking 
powder for alum and state results. Test them as follows : Put about 
a half gram of baking powder into a test-tube with 4 or 5 cc. of water 
and a little hydrochloric acid. Heat gently for a few moments, or 
until the solution is clear. If necessary, filter so as to obtain a per- 
fectly clear solution; then add a few drops of ammonia, or enough to 
make it alkaline. If alum is present, a more or less heavy precipitate, 
white or nearly transparent, will form. Some idea of the amount of 
alum present may be obtained by the quantity of the precipitate 
formed. This, however, should be verified by further tests. 

17. Clarifying Water. — In large cities obtaining their 
water supply from rivers which are muddy during cer- 
tain seasons of the year, considerable quantities of alum 
are used for settling the sediment. Weighed amounts of 
lime and alum are thrown into the settling basins ; the 
lime, on coming into contact with the water, is slaked, as 
you have learned, forming calcium hydroxide, thus : — 

CaO + H 2 = Ca(OH) 2 . 

This is a strong alkali, like ammonia, and forms in the 
water with the alum a precipitate of aluminum hydroxide, 
Al 2 (OH) 6 , just as the ammonia did with the alum in the 
baking powders. The reaction is as follows : — 

A1 2 (S0 4 ) 3 + 3 Ca(OH) 2 = Al 2 (OH) 6 + 3 CaS0 4 . 

Aluminum hydroxide is a gelatinous or starchy precipi- 
tate which in settling brings down with it practically all 
of the sediment, leaving the water clear. The only sub- 
stance added to the water is calcium sulphate, which, as 
seen in the study of calcium, simply renders the water a 
little more "hard." No trace of alum will be found to 
remain in the water, since the hydroxide is insoluble. 



ALUMINUM AND ITS COMPOUNDS 269 

18. Aluminum Hydroxide, Al 2 (OH) 6 . — This, as stated 
above, is a starchy, white, semi-translucent precipitate, 
formed when ammonia is added to any aluminum salt in 
solution. It may also be formed by adding caustic soda or 
potash to an aluminum solution, but in excess of these 
reagents it is soluble. 

Experiment 159. — Prepare some aluminum hydroxide as de- 
scribed above, using ammonia. Notice its appearance, and test its 
solubility in hydrochloric acid. What results? Write the reactions. 

Repeat the experiment, using caustic soda or potash instead of the 
ammonia. Add cautiously a drop at a time until the precipitate 
forms, and then an excess. State results. Write the two reactions, 
knowing that in the latter case potassium aluminate, K 3 A10 3 , a com- 
pound soluble in water, is formed. 

REVIEW OF WORK IN ALUMINUM 

Abundance of the metal. 

Form in which it occurs. 
Former methods of obtaining it and cost. 
Present methods of reduction. 
Description of the metal. 

Color, density, melting point, malleability, tenacity, conductivity, 

ease of oxidation, susceptibility of polish. 
Value of the metal in a practical way. 
Minor uses — Alloys. 
As a substitute for copper ; for iron. 
Compounds — Native. 

Difference between kaolin and clay. 

Uses of the former. 

Difference between emery and certain gems. 

Artificial compounds. 
Alums — Meaning of the term. 
Preparation of common alum. 
What is burnt alum. 
Uses of common alum. 

Chemistry of its use in clarifying water. 

Method of showing its presence in baking powders. 

Mordant — Meaning of term. 



CHAPTER XXIII 

TIN AND LEAD 
Tin: Sn = 118 

1. Source of Supply. — Almost the entire commercial 
supply of tin is obtained from the ore, cassiterite, Sn0 2 , 
sometimes called tin-stone. The ore probably received its 
technical name from an early appellation of the British 
Isles, Cassiterides. Extensive mines located at Cornwall, 
England, have been worked for hundreds of years. Long 
before the Christian era the Phoenicians brought back 
great quantities of ore from these mines ; yet even to 
this day they are very productive. They extend down 
into the earth thousands of feet, and far out, even under 
the bed of the ocean. The purest tin is said to be that 
obtained from India, known as Banca tin. Other sources 
of supply are Australia, Bolivia, and the Black Hills of 
Dakota ; the last-named mines, however, have not yet 
been well developed. 

2. Reduction of Ore. — With cassiterite there are usually 
found small quantities of arsenic in the form of arsenic 
sulphide, besides some other metals. After crushing the 
ore, it is strongly heated in a reverberatory furnace, where 
the volatile products, such as arsenic and sulphur, are 
expelled. At the same time the sulphides of the other 
metals are converted into sulphates, thus : — 

CuS + 2 2 = CuS0 4 . 
270 



TIN AND LEAD 271 

These sulphates are soluble in water, and to remove them 
the roasted ore is thoroughly washed; it is then mixed 
with coke or coal and reduced in a blast furnace, the 
carbon, as usual, serving to deoxidize the cassiterite. This 
last process may be indicated thus : — 

Sn0 2 + 2 C = Sn + 2 CO. 

Experiment 160. — If granulated tin is not to be had, some foil, 
procured from the tobacconist, will serve. Test the melting point of 
tin by holding a small piece in the flame. Try the effect of nitric acid 
upon tin. Also hydrochloric, and state results. Try it on charcoal 
with the oxidizing flame. State results. Does tin oxidize readily at 
ordinary temperatures ? 

Heat a piece of " tin plate " in the burner flame a moment, and then 
plunge it into cold water. Now rub the surface with a cloth mois- 
tened with diluted aqua regia, and wash with water. State the appear- 
ance of the surface. 

3. Characteristics of Tin. — Tin is a silvery white, lus- 
trous metal which does not tarnish in the air. It is some- 
what harder than lead, but melts at a lower temperature. 
By the oxidizing flame it may be converted into a white 
powder, Sn0 2 . It is highly crystalline in structure, as 
may be seen by removing the surface of a sheet or bar of 
tin with acid. The crystals may be prepared by melting 
some tin in a crucible, and when it is partially cooled, 
pouring out what is still molten. When a bar of tin is 
bent, these crystals rub together, and produce a distinctly 
audible sound, known as the " tin cry." This metal is 
very malleable, and may easily be beaten or rolled into 
thin sheets. It is soluble in aqua regia and in hydro- 
chloric acid, but ordinary nitric acid converts it into a 
white powder, metastannic acid, from which we may 
obtain stannic oxide, Sn0 2 , by expelling the water. 



272 MODERN CHEMISTRY 

4. Uses. — Because of its permanency in the air, tin is 
used in coating sheet iron, making what is known as tin 
plate. From this all "tin " vessels are made. In making 
sheet tin, the sheet iron is first thoroughly cleansed by 
immersion in dilute acids, then washed and dried. It is 
next dipped into a bath of molten tin, of which a thin 
coating adheres to the iron. A second and a third dipping 
increase the thickness of the coating. For outdoor work, 
such as roofing and guttering, a heavier quality of sheet 
iron is used, and the tin is generally alloyed with lead 
because the latter is much cheaper. 

5. In the form of foil, tin is extensively used for wrap- 
ping purposes ; at present, however, it is often adulter- 
ated with lead, especially in cases where the latter metal 
will be of no disadvantage. Tin foil amalgamated with 
mercury is also used frequently for the backs of mirrors. 

6. As stated elsewhere, tin is used extensively in alloys, 
among them being common solder, type metal, German 
silver, and the fusible metals. To these it imparts the 
property of a low melting point, that of the fusible metals 
being even lower than the boiling temperature of water. 
Spoons made from these metals, if dipped into a cup of 
hot tea or coffee, would rapidly melt and disappear. 

Compounds of Tin 

7. Stannous and Stannic Salts. — As is the case with 
many other metals, tin forms two general classes of salts : 
the stannous and the stannic. 

8. The Chlorides : Stannous, SnCl 2 , and Stannic, SnCl 4 . — 
The former is a white crystalline salt which may be pre- 
pared by dissolving tin in hydrochloric acid. It is a great 
reducing agent, and readily reduces mercury from certain 
of its salts. If a solution of stannous chloride be gradu- 



TIN AXD LEAD 273 

ally added to one of mercuric chloride, at first a white 
precipitate of mercurous chloride forms, and then by the 
addition of more of the tin solution, the precipitate slowly 
turns darker from the fact that the mercury is reduced to 
the metallic condition, though in a very finely divided 
form. The following reactions express the changes that 
take place : — 

2 HgCl 2 + SnCl 2 = Hg 2 Cl 2 + SnCl 4 ; 

Hg 2 Cl 2 + SnCl 2 = Hg 2 + SnCl 4 . 

Various other metallic salts are in a similar way reduced 
from the ic to the ous condition. The above reaction of 
stannous chloride and mercuric chloride upon each other 
may be used as a test for the presence of either metal. 

9. Stannic chloride may be prepared by dissolving tin 
in aqua regia. 

Experiment 101. — Dissolve some tin in hydrochloric acid, and 
boil nearly to dryness to expel the excess of acid. Take up with a 
few cubic centimeters of water, and gradually add it to a little of 
a solution of mercuric chloride in a test-tube. Do you obtain first a 
white precipitate, and afterward a gray one, becoming nearly black, 
as explained in the text above ? If necessary, warm gently. 

Dissolve a little tin in hydrochloric acid with a little nitric added; 
boil nearly dry, and after adding a few cubic centimeters of water, 
test with mercuric chloride, as before. State the results. 

10. The Sulphides : Stannous, SnS, and Stannic, SnS 2 . — 
These may be prepared by passing a current of hydrogen 
sulphide through solutions of stannous and stannic chloride, 
respectively. The former is a dark brown precipitate, 
the latter, bright yellow, closely resembling arsenic sul- 
phide. These two, though alike in some respects, differ, 
however, in that the latter is soluble in ammonium car- 
bonate, while the former is not. Stannic sulphide, more- 



274 MODERN CHEMISTRY 

over, is soluble in hot concentrated hydrochloric acid, 
while the arsenic is not. 

Experiment 162. — Add a drop or two of hydrochloric acid to a 
solution of stannous chloride, and pass through it a current of hydro- 
gen sulphide. Describe the precipitate which forms. Write the 
reaction. Treat a solution of stannic chloride in the same way ; what 
are the results ? Write the reaction. Take a part of this precipitate 
and add a little yellow ammonium sulphide; what happens? To the 
remainder add some ammonium carbonate solution. Is the sulphide 
dissolved ? 

11. Stannic Oxide, Sn0 2 . — This is principally of in- 
terest because it is the chief ore of tin. It is obtained 
when tin is strongly heated with the oxidizing flame ; it 
is pale yellow when hot, but white when cold. 

Exercise. — Write the reactions showing the preparation of all the 
above-named compounds. 

Experiment 163. — To determine whether any specimen of tin 
contains lead as an adulterant. Poisoning sometimes results from the 
canning of fruit in tin which is alloyed with lead. Procure any 
specimens of tin plate, or foil, and put upon them a drop or two of 
nitric acid. When dry, add a little of a solution of potassium iodide 
to the same spots. If the tin is adulterated with lead, bright yellow 
spots will appear, owing to the formation of Jead iodide. 

Lead: Pb = 207 

12. History. — Lead has been known from very early 
times because of the ease with which it is reduced from 
its ores. It is mentioned several times by biblical 
writers, but seems to have been confounded with tin. 
The two metals are spoken of b}^ Latin writers as black 
and white lead, respectively ; yet tin was the more expen- 
sive, and known to be suitable for soldering. 

13. Occurrence. : — The principal ore of lead is galena. 
or lead sulphide, PbS. It is a dark-colored, almost black, 



TIN AND LEAD 



275 



lustrous mineral, resembling somewhat metallic lead itself, 
but does not tarnish in the air as the metal does. It 
occurs in masses which tend to split up into cubical form ; 
it is widely distributed, but is usually found in what are 
called " pockets." It has very frequently associated with 
it ores of zinc, silver, iron, and some other metals. 

Experiment 164. — Pat into a small cavity in a stick of charcoal 
a little lead oxide, PbO, or minium, Pb 3 4 , or powdered galena, and 
heat with the reducing flame before the blowpipe. Do you obtain a 
metallic globule ? 

14. Reduction of the Ores. — This experiment illustrates 
in the main what is usually known as the " carbon re- 
duction." The furnace used in this method is not essen- 
tially different from the blast furnace shown under iron, 
for the reduction of iron ores. See the illustration on 
page 301. 

15. The Oxidation Process. — The second method may 
be called the " roasting " or " oxidation " process. The 
finely ground ore is placed upon the floor of the oven in a 
reverberatory furnace (see 
Fig. 57). The heat and 
flames are directed down- 
ward from the arching roof 
above upon the ore. In this 
way the upper layers are 
converted from the sulphide 
into the oxide, as seen in the 
reaction : — 




Fig. 57. 



PbS + 30 = PbO + S0 2 . 

The central portion of the mass being less strongly heated 
is converted into the sulphate, thus : — 



PbS + 2 On 



PbS0 4 . 



276 MODERN CHEMISTRY 

The bottom portions of ore, not receiving sufficient heat, 
undergo no chemical change. When sufficient time has 
elapsed to secure the above results, the strong draughts 
of air are shut off, in order to prevent the process of 
oxidation from going further ; thereupon the unchanged 
sulphide, PbS, reacts upon the oxide, PbO, and the sul- 
phate, PbS0 4 , already formed, whereupon metallic lead is 
obtained. We may represent this by the following reac- 
tions : — 

PbS + 2 PbO = 3 Pb + S0 2 , 

and PbS + PbS0 4 = 2 Pb + 2 S0 2 , 

or representing the complete change by a single reaction, 

2 PbS + 2 PbO + PbS0 4 = 5 Pb + 3 S0 2 . 

This is the process, probably, most commonly used, because 
the most economical, but it is adapted only to moderately 
rich ores. 

16. Reduction of Impure Ores. — When the lead ores 
are considerably mixed with the ores of other metals, the 
processes described above are not satisfactory. 

Experiment 165. — Suspend in a test-tube or bottle, about two- 
thirds full of a moderately strong solution of lead acetate, a strip of 
zinc. Allow it to stand for several hours without shaking. Notice 
the flaky crystals of lead that form on the zinc, giving what is called 
the "lead tree." After 24 to 48 hours carefully lift out the "tree," 
remove the crystals of lead, and notice how soft and porous the mass 
seems. Xotice how the zinc strip has changed. If you test the solu- 
tion in the bottle, you will find that it contains ziuc now instead of 
lead. That is, as the lead has slowly deposited upon the zinc, the 
zinc has likewise slowly dissolved. The chemical change is shown by 
the following reaction : — 

Pb(C 2 H 3 2 ) 2 + Zn = Zn(C 2 H 3 2 ) 2 + Pb. 



TIN AND LEAD 277 

17. For impure ores a similar method is adopted. They 
are mixed with scrap iron, and melted in a furnace, where- 
upon the lead, together with any silver present, is set free, 
and the iron combines with the other matters present. 

Thus : — 

PbS + Fe = FeS + Pb. 

18. Characteristics of Lead. — Many of these may be 
learned from the following simple experiments: — 

Experiment 166. — Take the globule of lead obtained in Experi- 
ment 164, and cut it with your knife. "What can you say of its hard- 
ness ? its color? its luster? Does it retain this luster? Test its melting 
point with the blowpipe and state results. What proof can you give 
that it oxidizes in the air? Try it on charcoal with the oxidizing 
flame ; what is seen upon the charcoal around the metal ? Compare 
it with silver in this respect. Put a small piece of lead into a test-tube 
and determine whether it is soluble in nitric acid ; in hydrochloric 
acid. 

19. Lead is a very heavy, soft, malleable, dark gray 
metal, with a specific gravity of 11.3 and a melting point 
of about 330°. It has a brilliant luster when first cut ; 
but, owing to the fact that it so readily oxidizes in the 
air, the surface is soon tarnished. This coating, however, 
protects the metal from further oxidation, and it is very 
durable. Lead is very different from iron in this respect, 
as the layer of rust, oxide, that forms upon the latter metal 
on exposure does not protect it. Lead is slightly soluble 
in ordinary water, and as all lead salts are very poisonous, 
water that has stood in leaden pipes any considerable 
length of time should not be used. The effects of lead 
compounds upon the human system are often seen among 
painters, who suffer from what is known as "lead colic." 

20. Uses for Lead. — This metal, as well as its com- 
pounds, has almost numberless uses. One of the most im- 



278 MODERN CHEMISTRY 

portant is for lead pipes in plumbing, used because they 
may be bent with ease. It is also employed largely for 
underground telephone conduits in cities, as well as for 
casings or sheaths for bundles of overhead wires. This 
pipe is made by forcing lead at a temperature near the 
point of solidification through an annular opening in a 
steel plate. 

21. Shot. — Another use of lead is in making shot and 
bullets. As stated elsewhere, for this purpose arsenic in 
small quantities is alloyed with the lead. One method is 
to allow the molten alloy to flow into a perforated vessel, 
from which the streams of metal fall long distances into 
water. In the descent the streams are broken into glob- 
ules, which before reaching the water have solidified. 
The various sizes and shapes thus obtained must next be 
sorted. The shot are allowed to roll down over inclined 
screens with a mesh of different sizes. The smaller shot 
will drop through first into one bin, the next size into a 
second, and so on. The irregular-shaped pieces will not 
roll through, and eventually make their way off the end 
of the plane. The shot are next polished by rotating in 
cylinders containing a little powdered graphite. 

22. Type Metal. — A third important use is in the 
manufacture of type for printers. This is made of an alloy 
of lead, tin, and antimony, or bismuth. The latter metals 
are used to give hardness to the alloy and to secure ex- 
pansibility at the moment of cooling. Owing to this 
property the type has clear, sharply cut faces, whereas 
lead alone would produce that having a battered or worn- 
out appearance. 

23. Solder. — A fourth use of lead is in solder, an alloy 
of tin and lead, the ordinary proportions being half and 
half. The tin is added to secure a low melting point, and 



TIN AND LEAD 279 

the proportions vary according to the use to which the 
solder is to be put. Lead is also used in making pewter, 
an alloy of lead and tin, and in storage batteries, but 
never for " lead " pencils. 

Compounds of Lead 

There are many of these, the most important among 
them being the following : — 

24. Lead Acetate, Pb(C 2 H 3 2 ) 2 , known also as Sugar of 
Lead, because of its sweet taste. It is a w r hite crystalline 
salt, which may be obtained by dissolving lead in vinegar 
or acetic acid, HC 2 H 3 2 . It is used frequently for dye- 
ing, and in medicine as an external application for ivy 
poisoning and in acute cases of erysipelas. 

25. Lead Chloride, PbCl 2 . — This may be prepared by 
adding hydrochloric acid to a lead solution, especially the 
acetate. It is a white solid, somewhat soluble in cold water, 
completely so in hot water, from which it crystallizes out 
upon cooling in small crystals that rapidly settle. 

Experiment 167. — To a few cubic centimeters of a solution of 
lead acetate, or nitrate, in a test-tube, add about 1 cc. of hydrochloric 
acid. Note the results. Write the reaction. Now add a little water 
and heat the contents of the tube to the boiling point. What hap- 
pens ? Allow it to cool and w^atch the tube meanwhile ; what happens? 
How do the two solids differ in appearance? 

26. Lead Sulphate, PbS0 4 . — This may be prepared by 
adding sulphuric acid to a soluble lead salt, as the acetate 
or nitrate. It is a heavy white salt, very slightly soluble 
in water and almost entirely insoluble in alcohol. 

27. Lead Nitrate, Pb(N0 3 ) 2 . — This salt is obtained when 
lead is dissolved in nitric acid. It is a white, crystalline 
compound, soluble in water. It is used somewhat in the 
laboratory. 



280 MODERN CHEMISTRY 

28. The Oxides. — There are several oxides of lead, the 
most important of which are PbO, litharge, or lead oxide ; 
Pb0 2 , lead peroxide ; and Pb 3 4 , minium, or red lead. 
This last is a deep red compound, used in plumbing to 
secure tight joints, and is sometimes regarded as a mixture 
of the other two, thus : — 

2 PbO + Pb0 2 = Pb 3 4 . 

It may be prepared from lead oxide, PbO, by heating. 
Litharge is a light brown-colored powder, obtained in 
large quantities when argentiferous lead ores are reduced 
by the cupellation process, and is always produced when 
lead is heated strongly in the air. It is used frequently 
in storage batteries, in preparing red lead, as stated above, 
and in making flint glass, to which it seems to impart the 
qualities of high refraction, almost perfect transparency, 
and softness. 

29. Lead Carbonate, PbC0 3 . — This is an insoluble white 
compound, which may be obtained by treating a solution 
of lead nitrate with one of ammonium carbonate. If 
sodium carbonate is used instead of the ammonium, a 
basic carbonate is obtained, or what may be represented 
by the formula, 2 PbC0 3 , Pb(OH) 2 , that is, two molecules 
of lead carbonate combined with one of lead hydroxide. 
In commerce this is known as white lead, and is used very 
extensively as paint. 

30. White lead is prepared by several methods, the 
oldest and perhaps the best being that known 
as the " Dutch method" (see Fig. 58). 
Glazed earthen jars, 8 or 9 inches in height, 
are used. About 3 inches from the bottom 
on the inside are some projections, upon 

Fig. 58. which a small board rests. Beneath the 




TIN AND LEAD 281 

shelf is vinegar, v, and above it a coil of sheet lead, q. 
Hundreds of these jars so prepared are placed side by 
side and covered with tan bark ; above them another 
layer of jars with a covering of bark is placed, and so on, 
to a considerable height. The whole is then buried under 
compost, which in decaying generates not a little heat. 
The fumes of acetic acid act upon the lead, gradually 
converting it into lead acetate. Then the carbon dioxide 
set free from the decaying tan bark combines with the 
acetate, slowly changing it into the basic carbonate. Sev- 
eral weeks are required for the completion of the process. 
The white lead is next removed from the jars, washed to 
dissolve out any lead acetate remaining, then ground in 
oil, and is ready for use 

31. Milner's Method. — Numerous attempts have been 
made to devise a method whereby white lead could be 
made quickly. One of these, which is fairly good, is 
Milners. Four parts of litharge, PbO, are mixed with 
one of common salt, Nad, and sixteen of water. The 
whole is ground together in a mill for 4 or 5 hours, 
then transferred to a leaden vessel, into which is con- 
ducted a current of carbon dioxide until the whole is 
neutral. 

32. White Lead by Electrolysis. — A current of elec- 
tricity is passed through a solution of sodium nitrate in 
water, in which a bar of lead is suspended. By the 
electric current the sodium nitrate is decomposed, forming 
caustic soda and nitric acid, thus : — 

NaN0 3 + H 2 = HX0 3 + XaOH. 

The nitric acid thus produced attacks the lead, and 
converts it into lead nitrate, Pb(N0 3 ) 2 . A second 
reaction follows, the lead nitrate and caustic soda 



282 MODERN CHEMISTRY 

combining to form lead hydroxide and sodium nitrate, 
thus : — 

Pb(N0 3 ) 2 + 2 NaOH = Pb(OH) 2 + 2 NaN0 3 . 

Thus we have produced again the same solution we had 
in the beginning, and only the lead needs to be renewed. 
The lead hydroxide thus obtained is treated next with so- 
dium carbonate, when basic lead carbonate is obtained, as 
follows : — 

2 Pb(OH) 2 + Na 2 C0 3 = 2 NaOH + PbC0 3 , Pb(OH) 2 . 

This process is continuous, very rapid, and is said to 
produce a fairly good quality of white lead. 

33. Lead Chromate, PbCr0 4 . — This is an insoluble com- 
pound of bright yellow color, and is easily prepared by add- 
ing a solution of potassium dichromate to one of a lead 
salt. It is used to a considerable extent as a paint, being 
sold under the name chrome yellow. 

Experiment 168. — Let the student prepare some of this pigment, 
and examine it. Use either potassium chromate or dichromate with 
a solution of lead nitrate or acetate. 

34. Lead Sulphide, PbS. — This is an insoluble com- 
pound, black in color, prepared by passing a current of 
hydrogen sulphide through a solution of lead nitrate or 
acetate. It has the same composition as native galena, 
but lacks the metallic luster. Galena is used in glazing 
pottery ware, except such as is to be used for articles of 
food. It is ground fine, mixed with pulverized clay and 
water, and the mixture washed over the pottery. When 
the vessels are strongly heated in ovens, the silica in the 
clay and the lead sulphide melt and form a glass which 
fills the pores of the clay. As such glazes are soluble, they 
are not suitable for pottery of all kinds. 



TIN AND LEAD 283 

Exercise. — Write the reactions expressing the preparation of all 
the lead salts described above. 

35. Identification. — Any solution of a lead salt may be 
identified by adding to it sulphuric acid or potassium 
dichromate, as in preparing the sulphate, or chromate, 
described above. A solution of potassium iodide is some- 
times used with the lead solution, and gives a bright 
yellow precipitate resembling chrome yellow. 

Experiment 169. — To determine the composition of common 
solder. Add to a small piece, not larger than a grain of wheat, 
about a cubic centimeter of concentrated nitric acid, and warm 
gently. When the alloy has disappeared, and the white powder 
which has formed is settled, decant the clear solution into an evapo- 
rating dish. Add some water to the white powder, and decant again. 
Now evaporate the solution decanted nearly to dryness, add a little 
water, and make two tests for lead with separate portions of it, 
according to method of identification suggested above. State your 
conclusion. 

To the white powder obtained at the beginning, add a little strong 
hydrochloric acid, and heat until solution is secured. Boil down 
nearly to dryness, add 25 to 50 cc. of water, and through part of it 
pass hydrogen sulphide ; to another portion add slowly, drop by drop, 
mercuric chloride. State results. From these can you determine 
what metal you have ? See section 9, page 273, and compare results. 

SUMMARY OF CHAPTER 

History of tin and lead. 

Occurrence of each — Chief ore, and its composition. 

Principal tin mines — Description. 
Reduction of the ores. 

Wherein similar. 

Purpose of the roasting in each case. 

Description of a second method of reducing lead. 
Furnace used. 
Chemical changes and reactions. 

Experiment of " lead tree " — Description — Purpose, 



284 MODERN CHEMISTRY 

Characteristics of lead and tin — Compare them in 
Color. 
Density. 
Hardness. 
Melting point. 
Malleability. 
Tendency to oxidize. 
Tendency to crystallize. 
Solubility in acids. 
Uses. 

Sheet tin — What is it ? — Its use — How made ? — Why ? 
Alloys of tin — Properties secured by the tin. 
Foil — Purposes. 
Lead pipes — Use — How made ? 
Shot — Manufacture of. 
Type metal. 
Solder. 
Compounds. 

The oxides of tin and lead — Compare them. 

Uses and preparation. 
Stannous chloride ; lead chloride. 
Preparation of each. 
Interesting facts about each. 
The Sulphides — Preparation of each. 
Appearance of each. 
Uses of PbS. 
Other important lead compounds. 

Sugar of Lead — Chemical name and formula. 
How prepared. 
Uses. 
White lead — Composition. 

Best way of preparing; give plan and chemical reac- 
tions. 
Electrolytic method. 
Chrome yellow. 

How prepared in laboratory. 
Appearance and uses. 
Usual method of identification of lead and tin salts. 
Analysis of common solder. 



CHAPTER XXIV 

ARSENIC, ANTIMONY, BISMUTH 
Arsexic : As = 75 

1. Source of Supply. — In limited quantities metallic 
arsenic is found free in one or two countries of Europe, 
especially Germany. The greater part, however, is ob- 
tained from arsenical pyrite, that is, iron pyrite, FeS 2 , in 
which arsenic has replaced an atom of sulphur, thus, 
FeAsS. It also occurs in combination with other metals, 
such as zinc and nickel, and with sulphur, as red arsenic 
sulphide or realgar, As 2 S 2 , and yellow arsenic sulphide, 
As 2 S 3 , or orpiment. 

2. Reduction of the Ores. — As already stated, arsenical 
pyrite is most commonly used for the production of arsenic. 
This ore is first roasted in ovens at a moderately strong 
heat, by which it is oxidized, thus : — 

2 FeAsS + 5 2 = Fe 2 3 + As 2 3 + 2 S0 2 . 

The last two of these products are volatile and are passed 
over into cold chambers, where the oxide of arsenic con- 
denses in a white sublimate. This is next mixed with 
powdered charcoal, put into retorts, and heated. The 
arsenic oxide is deoxidized by the charcoal, the metallic 
arsenic vaporizes and is condensed in cold chambers, 
thus : — 

As 2 3 + 3C = 2As + 3CO. 

285 



286 MODERN CHEMISTRY 

Experiment 170. — Mix well a little arsenic trioxide and some 
powdered charcoal and put into one end of a piece of hard glass 
tubing. Now, heat strongly in the Bunsen flame, holding the tube in 
a slanting position, with the cooler end up. Notice the deposit form- 
ing. Describe it. What has been the effect of the charcoal ? 

This experiment illustrates the commercial method of preparing 
arsenic. 



3. Characteristics of Arsenic. — In many respects arsenic 
is not unlike some of the non-metallic elements, notably 
phosphorus. It forms compounds with hydrogen and 
oxygen similar to those of phosphorus, and with several 
metals a variety of salts in which arsenic is the acid- 
forming element ; for example, sodium arsenate, Na 3 As0 4 , 
nickel arsenide, NiAs, etc. In general appearance, how- 
ever, arsenic is more like the metals. Thus, it is of a 
dark gray color, with metallic luster when freshly broken, 
has a marked tendency to crystallize, and tarnishes slowly 
in moist air. It vaporizes without melting, and when in 
the form of vapor has a sickening garlic odor. It is of 
medium density, and, like phosphorus, has four atoms to 
the molecule. It is but little acted upon by nitric or 
hydrochloric acid, but dissolves readily in aqua regia or 
nascent chlorine. It has strong affinity for chlorine, and 
if it be finely powdered and sifted into a bottle of the gas 
it burns readily. 

Experiment 171. — Examine some crystals of metallic arsenic and 
notice their color and general appearance. Are they malleable ? Heat 
a small piece on charcoal with the blowpipe. Notice the odor. Does 
the arsenic melt V What becomes of it ? 

4. Uses. — In the metallic form arsenic has little use 
except in making shot. With lead it forms an alloy that 
is considerably harder than the former metal, and at the 
same time one which, in the molten condition, is much 






ARSENIC, ANTIMONY, BISMUTH 287 

more mobile. This property of the arsenic alloy is of 
value in the manufacture of shot by the ordinary method, 
for the shot made from it are more perfect in shape than 
those made from a metal more viscous, like lead. 

Compounds of Arsenic 

Experiment 172. — To study the characteristics of arsine, AsH 3 . 
Prepare a flask for the generation of hydrogen from zinc and sulphuric 
acid, as on page 39, and attach a jet. After a few moments, or 
when sufficient time has elapsed for the air to be expelled, wrap a 
towel around the flask or inclose in a small box 
with an opening through the cover, as seen in the 
figure, and light the jet. You have hydrogen 
burning. Hold a cold porcelain dish against the 
flame and notice that no deposit forms upon the 
dish. 

Xow, add to the hydrogen flask a little arsenic 
trioxide, dissolved in dilute hydrochloric acid. 
Again light the jet, and notice how the color of 
the flame has changed. Hold a cold dish against G * 

the jet as before. Is any deposit formed? What 

that you have already seen does it closely resemble? The gas being 
generated is arsine. Xow, hold a beaker or test-tube over the burning 
jet, and notice whether there are not two different deposits formed. 
Can you decide what they are ? Write the reaction that takes place 
when arsine burns. 

5. Arsine. — Arsine, AsH 3 , is also known as arseniu- 
reted hydrogen, or hydrogen arsenide. It is a compound 
of considerable interest, because it is always prepared in 
testing for arsenic in cases of suspected poisoning. The 
method used is the one described in the experiment 
above. This is known as Marsh's test, and is so exceed- 
ingly delicate that mere traces of arsenic, even so low as 
one part in several hundred thousand, can be detected. 
Care should be taken, however, to see that the zinc is 




288 MODERN CHEMISTRY 

perfectly free from arsenic. Antimony gives a spot con- 
siderably like that of arsenic seen above, but the latter 
may be detected by treating with a solution of bleaching 
powder, in which the arsenic spots are soluble, while the 
others are not. 

6. Let us study the reactions that take place. First, 
by the reaction of sulphuric acid and zinc upon each other 
hydrogen is produced, thus : — 

Zn + H 2 S0 4 = H 2 + ZnS0 4 . 

The hydrogen atoms in the nascent condition, instead of 
uniting with one another to form molecules of hydrogen, 
unite with the arsenic present, forming hydrogen arsenide, 
AsH 3 . This may be represented thus : — 

AsCl 3 + 6 H = AsH 3 + 3 HC1. 

7. Characteristics of Arsine. — This is a colorless, ex- 
ceedingly poisonous gas, which burns with a pale violet 
flame, giving off white fumes of the trioxide As 2 3 . 

2 AsH 3 + 3 2 = As 2 3 + 3 H 2 0. 

Both of these products may be seen if a cold beaker or 
test-tube be held over the burning jet of arsine. If a cold 
dish is held against the flame, the temperature is lowered 
below that required for the combustion of arsenic, and it 
is therefore deposited in the metallic form, while the 
hydrogen continues to burn. What does the experiment 
teach regarding the kindling point of hydrogen ? 

8. The Oxides of Arsenic. — Corresponding to the two 
oxides of phosphorus we have two of arsenic, the trioxide, 
As 2 3 , and pentoxide, As 2 5 . Only the former is of im- 
portance. It occurs in two or three forms, the white 
powder being the most common. It is usually sold under 



ARSENIC, ANTIMONY, BISMUTH 289 

the name " arsenic " or white arsenic, but is also called 
arsenious acid. It has a sweetish taste, is slightly soluble 
in cold water, more so in hot, in hydrochloric acid, and in 
caustic soda. It is very poisonous, but acts somewhat 
slowly. An antidote for it is .ferric hydroxide, prepared 
by treating a ferric salt in solution with ammonia ; the 
precipitate must be filtered out and washed. Magnesia, 
MgO, is also suggested, and is used more often because it 
is to be had already prepared. 

9. Arsenic trioxide is used by taxidermists in curing 
the skins of animals ; it is an ingredient of many poisons, 
but is also often prescribed by physicians as a blood purifier, 
especially for removing facial eruptions. It is thought to 
beautify the complexion, and has a tendency to produce 
fat. Because of the latter property it is sometimes fed to 
old horses to prepare them for the market. It stimulates 
the action of the heart and renders breathing easier ; on 
this account it is said to be used by some mountain climb- 
ers. These apparent benefits are but temporary, however, 
and a discontinuance of its use is attended by all the 
symptoms of serious arsenic poisoning. 

Experiment 173. — Examine a sample of arsenic trioxide and note 
its general appearances. Test its solubility in diluted hydrochloric 
acid, also in caustic soda. Which is the better solvent? Use only 
small quantities of the trioxide. Save the solution. 

10. Paris Green ; Scheele's Green. — This is a very 
poisonous, bright green powder, used often for coloring 
and tinting and as an insect exterminator. 

Experiment 174. — Let the student prepare this compound, thus: 
To a few cubic centimeters of a solution of copper sulphate in a 
test-tube add ammonia, drop by drop, until the precipitate which 
forms at first just dissolves. Now, add gradually a solution of arsenic ; 
a bright green precipitate will form. If too blue, not enough arsenic 



290 



MODERN CHEMISTRY 



has been added. This is one of the easiest methods of detecting 
arsenic if present in considerable quantities. It is known as Scheele's 
test. 

11. Arsenic Trisulphide, As 2 S 3 . — This is a bright yel- 
low powder obtained by passing a current of hydrogen 
sulphide through a solution of arsenic. It is soluble 
in ammonium carbonate, which distinguishes it from a 
similar compound of tin, SnS 2 , also yellow. It is also 
soluble in yellow ammonium sulphide, but not in hydro- 
chloric acid. 

Experiment 175. — Let the student prepare this compound by 
passing hydrogen sulphide through a solution of arsenic trioxide in 
water acidulated with hydrochloric acid. Divide the yellow precipi- 
tate into two or three parts and test its solubility in hydrochloric acid 
and in ammonium sulphide and carbonate. 

Antimony : Sb = 120 

12. Source of Antimony. — This element is found free in 
very small quantities only, but frequently occurs with the 
ores of other metals, such as lead, copper, and iron. Its 
principal ore is stibnite, Sb 2 S 3 , and from this the commer- 
cial supply is obtained. 

13. Reduction of the Ore. — There are two methods used 
for reducing antimony ores. The first consists in heating 

the sulphide in a reverberatory 
furnace, whereby the ore is re- 
duced to an oxide, thus : — 

Sb 2 S 3 + 5 2 = Sb 2 4 + 3 S0 2 . 

Then the tetroxide, thus formed, 
is mixed with charcoal, and again 
heated in a furnace, when metallic antimony is obtained, 
thus : — 

Sb 2 4 + 4 C = 4 CO + 2 Sb. 




Fig. 60. 



ARSENIC, ANTIMONY, BISMUTH 291 

Experiment 176. — In a cavity in a piece of charcoal place a little 
antimony tartrate, mixed with sodium carbonate, and moisten with 
a few drops of water. Xow heat strongly with the reducing flame. 
What do you obtain ? Preserve for the next experiment. 

14. This illustrates the method of reduction described 
above, and, it will be noticed, is in accord with the general 
plan of reducing metallic ores, — first reducing them to 
the form of an oxide by roasting them, and then deoxi- 
dizing them, by heating with carbon. 

15. Another Method. — This consists in mixing the ore, 
antimony sulphide, with iron, and melting the whole in a 
furnace. The iron combines with the sulphur, and pre- 
cipitates the antimony, thus : — 

Sb 2 S 3 + 3 Fe = 3 FeS + 2 Sb. 

16. Characteristics of Antimony. — Owing to the fact 
that, like phosphorus, nitrogen, and other non-metallic ele- 
ments, antimony forms oxides which are the anhydrides of 
acids, it is sometimes regarded as a non-metallic element. 
It is, however, of a highly lustrous metallic appearance, 
steel-gray in color, notably crystalline in structure, heavy, 
and so very brittle that it is easily reduced to a powder. 

17. Antimony combines energetically with chlorine, 
bromine, and iodine, in contact with all of which, when 
finely powdered, it quickly takes fire. Upon bromine, 
sufficient heat is generated to melt the antimony, and it 
spins around as does sodium upon water, burning all the 
time. At ordinary temperatures, the metal does not 
readily tarnish in the air, but by means of the oxidizing 
blowpipe flame it is converted into a white oxide, Sb 2 3 . 
It is only slightly acted upon by hydrochloric acid, but 
nitric acid converts it into a white powder, and it is 
readily soluble in aqua regia, forming antimony chloride, 



292 MODERN CHEMISTRY 

SbCl 3 . One of its most valuable properties is that of 
expanding somewhat upon cooling. 

Experiment 177. — To illustrate some of the above-mentioned 
properties. Take the metallic bead obtained in the preceding experi- 
ment, and learn whether it is magnetic. Test it with a hammer on 
an anvil to learn whether it is malleable. Notice its color and appear- 
ance. Put a portion of it on charcoal and try the oxidizing flame. 
What are the results ? How does it differ from arsenic treated thus ? 
Test the solubility of the metal in nitric acid ; in aqua regia. State 
results in each case. Boil nearly to dryness the latter solution, and 
add water. What happens? Treat this with tartaric acid, and state 
results. 

18. Uses. — Because of its property of expanding when 
it solidifies, antimony is used very extensively in mak- 
ing type metal, britannia ware, and other similar alloys. 
Antimony may be obtained in a powdered or amorphous 
condition by immersing a strip of zinc in a solution of 
some antimony salt, as the chloride or tartrate. The 
principle underlying is the same as that in the second 
method of reducing the ore, described already. This 
antimony black, as it is called, is a dark-colored, finely 
divided powder, and is sometimes used in giving plaster 
figures a metallic appearance. 

Compounds of Antimony 

19. There was a time when the compounds of antimony 
were extensively employed in medicine, but owing to their 
exceedingly poisonous character, their use was prohibited 
by law, and their applications now are considerably 
limited. 

20. Stibine, Atimoniureted Hydrogen, SbH 3 . — This 
gas, known also as hydrogen antimonide, corresponding to 
similar compounds of arsenic and phosphorus, is usually 



ARSENIC, ANTIMONY, BISMUTH 293 

prepared from nascent hydrogen and some antimony com- 
pound, just as arsine was prepared in Experiment 172. 
It is a combustible gas, which burns with a green flame, 
and deposits upon a cold dish held against this flame a 
black spot resembling that of arsenic, but not so lustrous. 
It is also less volatile if heated, and is insoluble in a solu- 
tion of calcium or sodium hypochlorite. 

Experiment 178. — Prepare stibine exactly as you did the arsine, 
using the same precautions. Test the spots with a solution of bleach- 
ing powder or sodium hypochlorite, and verify the statements made 
above. 

21. Oxides of Antimony. — None of the three oxides of 
antimony is of any importance. The trioxide, Sb 2 3 , and 
pentoxide, Sb 2 5 , are the anhydrides of the acids, antimo- 
nous and antimonic, corresponding to those of nitrogen 
from the similar oxides. 

Sb 2 3 + 3 H 2 = 2 H 3 Sb0 3 . 

Sb 2 5 + 3 H 2 = 2 H 3 Sb0 4 . 

22. The Chlorides of Antimony. — When antimony is 
dissolved in aqua regia, as in Experiment 177, above, and 
the solution evaporated, antimony trichloride, SbCl 3 , a 
white crystalline salt, is obtained. It was formerly known 
as " butter of antimony," from the thick oily appearance 
which it assumes before solidifying. Upon adding water 
to this compound, a white precipitate is formed, which is 
known as basic antimony chloride, or antimony oxychlo- 
ride, SbOCl. The reaction may be expressed thus : — 

SbCl 3 + H 2 = SbOCl + 2 HC1. 

The trichloride has given up two atoms of its chlorine, 
and has taken in their place one atom of bivalent oxygen. 



294 MODERN CHEMISTRY 



urx,- feW ™«' 



© (a) @><g) 



This oxychloride is soluble in tartaric acid, but not in 
water. 

23. Antimony Trisulphide, Sb 2 S 3 . — This is obtained arti- 
ficially by passing a current of hydrogen sulphide through 
an antimony solution. It is of a beautiful orange color, 
soluble in yellow ammonium sulphide, and also in strong 
hydrochloric acid. 

Bismuth: Bi = 208 

24. Source of Supply. — Most of the commercial supply 
of bismuth is obtained from Saxony. It is usually found 
free, but alloyed with small quantities of several other 
metals. It also occurs in two ores : the sulphide, Bi 2 S 3 , 
known as bismuthite, and the oxide, Bi 2 3 . 

25. Reduction. — When obtained from native bismuth, 
as it usually is, the process consists of little more than 
simply heating to melt the bismuth ; the other metals 
found with it have a higher melting point, and remain 
unchanged. In the case of the ores, if bismuthite is used, 
it is treated as the sulphides of other metals are, first con- 
verted into an oxide, and then heated with charcoal. Let 
the student write the reactions representing the two 
steps. 

26. Characteristics. — Like antimony, bismuth is a hard, 
brittle, distinctly crystalline metal. It is steel-gray in 
color, having somewhat of a golden reflection, or upon 
some surfaces a purplish hue. It has a low melting point, 
being just above tin in this respect, expands upon solidi- 



ABSENIC, ANTIMONY, BISMUTH 295 

fying, and is permanent in the air at ordinary tempera- 
tures ; at a red heat it oxidizes to a light yellow powder. 
It unites readily with bromine and chlorine, and if sifted 
into them takes tire at once. 

Experiment 179. — If no bismuth is to be had in the laboratory, 
prepare a little by heating bismuth nitrate, mixed with sodium carbo- 
nate and moistened, on charcoal with the reducing flame. 

Note the color of the metallic bead ; test its hardness and mallea- 
bility, and learn whether it is magnetic. Dissolve a portion of the 
bead obtained in nitric acid, boil nearly dry, and add water. What 
forms? Treat with tartaric acid in solution. Compare results with 
similar tests with antimony. How do they differ? 

27. Uses. — In the metallic form bismuth has but little 
use, except in alloys. To these it imparts the properties 
of low fusing points and of expansibility. For these 
reasons it is used in stereotyping, and for similar purposes 
where clearly defined copies are demanded. Bismuth is 
also used for making safety plugs in boilers, and for very 
fusible alloys, such as Wood's alloy, which melts at about 
60° C. 

Experiment 180. — Put into an iron spoon about 2 g. of bismuth, 
1 g. of lead, and 1 g. of tin, and melt them. When cold put into a 
beaker of boiling water. What happens ? 

28. Most of the bismuth produced at the smelters is 
converted into its compounds and used in a medicinal way. 

Compounds of Bismuth 

29. Two Classes of Compounds. — Like antimony, bis- 
muth forms two classes of compounds : the ordinary, and 
the basic or bismuthyl. These are best seen in the nitrates, 
Bi(N0 3 ) 3 , in which the bismuth atom has its true valence, 
three, and in the basic nitrate, BiON0 3 , in which one 



296 MODERN CHEMISTRY 

atom of oxygen has replaced two of the groups of N0 3 . 
It may be graphically shown as follows : — 



9 



Ordinary. Basic (Bismuthyl) . 

30. The first of these is a white crystalline salt, which 
is prepared by dissolving metallic bismuth in nitric acid„ 
It has little use, except in the preparation of other com- 
pounds of bismuth. The basic nitrate, sold at drug-stores 
as the subnitrate, or simply as " bismuth," is a white 
powder, obtained from the ordinary nitrate by the addi- 
tion of water, whereupon a fine white precipitate falls, 
thus : — 

Bi(N0 3 ) 3 + H 2 = BiON0 3 + 2 HN0 3 , 

or more properly, considering the water of crystallization, 

Bi(N0 3 ) 3 , 2 H 2 + H 2 0-BiON0 3 , H 2 + 2 HN0 8 + H 2 0. 

This is used largely as a cosmetic, and for relieving the 
irritation of chafed or chapped skin ; also in cholera and 
kindred diseases, and in acute dyspepsia. 

31. Bismuth Trioxide, Bi 2 3 . — This is also called bis- 
muth ocher, the chief ore of bismuth, but may be obtained 
artificially by heating the metal in the oxidizing flame. 
It is of a deep yellow color when hot, but yellowish white 
when cold. Its principal use is as a paint. 

32. Bismuth Trichloride, BiCl 3 . — This may be prepared 
by heating bismuth in chlorine gas. If water is added 
to it, the basic bismuth chloride, or oxychloride, BiOCl, 
is formed, as is the case with antimony. The latter, how- 



ARSENIC, ANTIMONY, BISMUTH 



297 



ever, is soluble in sodium tartrate or tartaric acid, but the 
former is not. Basic bismuth chloride is a fine white 
powder, and is used as a paint, known as " pearl white." 
33. The Nitrogen Group. — From the similarity of their 
compounds, and their chemical affinity, nitrogen, phos- 
phorus, arsenic, antimony, and bismuth are often classed 
together and called the nitrogen group. The following 
table will give a comparative view of their more impor- 
tant compounds : — 



N=14 

Nitrogen 


P = 31 

Phosphorus 


As = 75 
Arsenic 


Sb = 120 
Antimony 


Bi = 208 
Bismuth 


NH 3 
NA 


PH 3 

P 2 3 

PA 


AsH 3 

As 2 3 
As 2 5 

ASC1 3 


SbH 3 
Sb 2 3 

Sb 2 O s 
SbCl 3 

SbOCl 




BiA 
BiA 
BiCl 3 
BiOCl 













SUMMARY OF CHAPTER 

Comparative Study of Arsenic, Antimony, and Bismuth. 
Sources of the metals. 

Wherein alike. Wherein different. 

Reduction of the ores. 

Wherein similar — How similar to reduction of other metallic 

ores. 
In what respects different. 

Description of experiments illustrating methods. 
Characteristics of the group. 

Compare two of them with the non-metals. 
Wherein are they all metallic in character. 
Compare in 

Color. Melting point. 

Density. Tendency to oxidize. 

Hardness. Solubility in acids. 

Malleability. 



298 MODERN CHEMISTRY 

State any special characteristics, hot common. 

Compare bismuth and antimony as to certain classes of salts 

formed by each. 
Compare arsenic and antimony in the same way. 
Uses of each. 

Special use for metallic arsenic — Reason. 
Same for antimony, and reason. 

Antimony black — How made ? — Use ? 
Same for bismuth, and reason. 
Compounds. 

Compare the hydrogen compounds of arsenic and antimony as to 
Method of preparing and chemical action. 
Characteristics of each. 
How distinguish one from the other ? 
Products formed when AsH 3 burns. 
Experimental proof. 
Oxygen compounds. 

Names and formulae. 
Important one of arsenic — Why ? 
Appearance and uses. 

Physiological action — Compare with antimony. 
Antidotes. Solvents. 

Appearance and use of bismuth oxide. 
Sulphides. 

Names and formulae. 
Method of preparing. 
Appearance of each. 
How distinguish As 2 S 3 from SnS 2 ? 
How distinguish As 2 S 3 from Sb 2 S 3 ? 

How distinguish an antimony salt from one of bismuth ? 
Special compounds. 

Paris green — Experimental preparation. 

Appearance — Uses. 
Butter of antimony — Chemical name and formula. 
Means of identifying. 

For arsenic — Marsh's test ; Scheele's test. 
For bismuth and antimony. 
Comparison of the nitrogen group. 

Compounds with hydrogen, oxygen, chlorine, etc. 



CHAPTER XXV 

IRON, NICKEL, COBALT 
Iron : Fe = 56 

1. Distribution. — Iron, the most useful of all metals, 
is also the most abundant and most widely distributed. 
It is found in nearly all clays and soils, and from these 
is taken up by plants, and through them makes its way 
into the animal economy. The color of many soils, rocks, 
and minerals is due to the presence of iron in some form. 
Pure iron does not occur in any considerable quantities, 
except in meteorites, of which some weigh many tons. 
The largest meteorites ever found were discovered by 
Lieutenant Peary in his Arctic explorations. One of 
these, weighing nearly one hundred tons, was brought 
back and placed in the Brooklyn Navy Yard. Meteorites 
are found to consist of iron, about 93 per cent, and nickel, 
7 per cent. 

2. Iron Ores. — A large number of iron ores are known, 
among which are the following : — 

3. Magnetite, Fe 3 4 . -. — This is also known as lode-stone, 
on account of its magnetic properties. 

4. Hematite, Fe 2 3 . — This ore received its name from 
the Greek word for blood, because of the red streak it 
gives 'on porcelain. It is a very abundant ore : two 
knobs, Iron Mountain and Pilot Knob, of the Ozark Range 
in Missouri, consist almost entirely of hematite, in masses 
ranging all the way from " the size of a pigeon's egg to 
that of a medium-sized church." 

299 



300 



MODERN CHEMISTRY 



5. Limonite, Fe 2 3 , 2 H 2 0. — An ore resembling hema- 
tite, which gives a yellow streak on porcelain. 

6. Siderite, FeC0 3 ; Spathic Iron Ore. — This is common 
in some localities, has a gray to brownish red color, and 
often contains manganese. 

7. Iron Pyrites, FeS 2 , is a very abundant ore, but on 
account of the difficulty experienced in reducing it, it 
is not used, except for the manufacture of sulphuric acid. 
It is commonly known as ." fool's gold." 

8. Reduction of Ores. — This is accomplished in a blast 
furnace, the essential features of 
which are shown in Figure 61. The 
furnace is from 50 to 75 feet in 
height, supported by masonry, and 
strengthened with boiler plate. It 
is lined inside with fire-brick. Near 
the bottom some pipes, PP, enter the 
furnace. These are called tuyeres, or 
blast-pipes, and furnish a powerful 
blast of air. Just below these is an 
opening, #, where the slag is drawn 
off, and below this another opening, 
i", for drawing off the iron. 

9. The furnace is charged from the top ; first, wood for 
kindling being placed in the bottom, then alternate layers 
of coke and iron ore mixed with limestone. These are all 
dumped upon the cone-shaped top, (7, which fits air tight 
and works automatically. When the ore and other mate- 
rials fall upon the top, it lowers mechanically and allows 
the charge to roll into the furnace. Many of the gases 
formed in the interior of the mass are combustible, and 
are conducted off through the pipe M, and are burned in 
other furnaces. 




Fig. 61. 



IBON, NICKEL, COBALT 



301 



Figure 62 shows a perspective view of the blast furnace. 
The various materials are lifted to the top by elevators ; 




Fig. 62. — Perspective View of Blast Furnace. 

the molten iron is drawn off and molded in trenches in 
the ground under the shed. 



302 



MODERN CHEMISTRY 



10. In accordance with the usual method of reducing 
metallic ores, the oxides of iron are mixed with coke and 
limestone and strongly heated. If an ore, not an oxide, 
is used, it is first calcined to convert it into an oxide. 
The coke serves as a deoxidizing agent, and the lime- 
stone, used as a flux, is melted, and renders the iron ore 
more fusible. The limestone then combines with the 
silica always present in the ore, and forms a molten glass, 
or slag, which floats upon the iron and prevents its oxida- 
tion by the strong currents of air. The iron thus obtained, 
nevertheless, absorbs in the intense heat of the furnace 
considerable quantities of sulphur, phosphorus, carbon, 
and silica, and in this impure form is not suited to many 
of the numerous demands for iron. From the blast fur- 
nace it is run off through trenches into molds, 2 to 4 feet 
long, which are called "pigs," and the iron is known as 
cast or pig iron. It is very brittle, coarse grained, and 
contains from 5 to 10 per cent of impurities. 

11. Wrought Iron. — This variety is prepared in what 
are known &s puddling furnaces. In these the low arch- 
ing roof deflects the flames down- 
ward upon the broken cast iron. 
The furnace is lined with ferric 
oxide, Fe 2 3 , and as the cast 
iron melts, the carbon which it 
contains combines with the oxy- 
gen in the lining. By stirring 
the molten mass, or puddling, 
as it is called, the whole is 
gradually purified, until finally, 

as it is much more difficult to melt pure iron, the whole 
mass becomes pasty. This pasty mass, bloom, as it is 
called, is then removed and hammered with trip-hammers, 




Fig. 63. 



IB OX, NICKEL, COBALT 



303 



a process which drives out any remaining slag, and renders 
the iron malleable. 

12. Steel. — Formerly steel was made from wrought 
iron by embedding bars of the latter in finely powdered 
charcoal and keeping at a red heat for about ten days. 
During this time the bars of iron slowly absorbed more 
or less carbon, and were converted into steel. Besides 
the expense and the length of time required in this 
process, there were two other serious objections to it : 
first, that there was no possible way of controlling accu- 
rately the amount of carbon taken up by the iron, and 
second, a steel bar was obtained which was not at all 
uniform in quality and texture. 

13. Present Method of Manufacture. — At present, steel 
is made directly from cast iron, by the Bessemer process. 
An egg-shaped vessel, called a converter, is used. It is 
securely bound with boiler iron, 
and is lined with ganister, a 
siliceous earth, fusible only at a 
very high temperature. The 
converter, which will hold ten 
or more tons of iron, is mounted 
on axes, or trunnions. One of 
these, J., in Fig. 64, is hollow, 
so that a blast of air may be 
forced through it when the 
converter is in a vertical posi- 
tion. This trunnion opens into 
a pipe, P, which extends down the outside of the converter 
and opens into the tuyere box, B, beneath the body of the 
converter. Through the tuyere box, numerous small 
openings admit the air to the converter and its con- 
tents. 




Fig. 64. A Converter. 



304 MODERN CHEMISTRY 

14. Bessemer Process. — About ten tons of pig iron are 
placed in a cupola furnace, that is, one resembling a blast 
furnace in most of its details, but considerably smaller. 
When the iron is melted, it is run into the converter. 
Immediately the blast of air is turned on, and, bubbling 
up through the molten iron, the oxygen unites with the 
carbon and other impurities, burning them out. No heat 
is used in the operation except what is evolved by the 
combustion of the impurities themselves. About twenty- 
five or thirty minutes are required for the completion of 
the operation, during most of which time a brilliant 
shower of sparks is thrown from the mouth of the con- 
verter. This is represented in colors by the frontispiece. 
The converter on the right is shown in action ; the one 
on the left, at the close of the process, discharging the 
molten steel into a pot, from which it will be poured 
into molds. 

15. When the mass of flame and sparks no longer 
issues from the converter, the workmen know that the 
cast iron has had its impurities entirely removed, and is 
now wrought iron. Next, a weighed quantity of spiegeleisen 
or manganese iron, containing a known amount of carbon, 
is thrown into the converter, and in a moment or two the 
process is complete. In this way, in thirty minutes or 
less, ten tons of steel are obtained at a cost only a fraction 
of what it would be by former methods. 

16. Basic-lining Process. — If the iron ore contains 
much phosphorus, the converter is lined with lime- 
stone, which, during the process of oxidation, takes up 
the phosphorus from the iron, and is converted into 
calcium phosphate. This is known as the basic-lining 
process and was put into practical use by the inventors, 
Thomas and Gilchrist. 



IRON, NICKEL, COBALT 305 

17. Tempering Steel. — Tempering consists in harden- 
ing steel, by heating and then suddenly plunging into 
cold water or oil. Tempered in this way, it becomes 
much less malleable, but can take and hold a sharp edge. 
Different instruments require steel that has been heated 
to different temperatures ; thus, surgical instruments after 
being hardened are again heated to about 225° or till a 
yellow film of oxide appears upon the surface. For 
ordinary cutlery, a temperature of about 250° is used, 
indicated by the appearance of a brown film, while clock 
and watch springs and such forms as require great elas- 
ticity are made of steel heated till blue, or about 290°. 
By heating any form of steel strongly and then cooling 
very slowly, the temper is " drawn," or removed, and the 
metal becomes like ordinary wrought iron. 

18. Comparison of the Three Forms. — 

Cast Iron. Steel. Wrought Iron. 

Impurities— 5 to 10%. 1 to 2 %. 0.36 to 0.5 %. 

Brittle. Somewhat malleable. Very malleable. 

Coarse grained. Fine grained. Very fine grained. 

Cannot be tempered. May be tempered. Cannot be tempered. 

Lowest melting point. Medium melting point. Highest melting point. 

19. Uses of Iron. — This is preeminently the " Steel 
Age." Day by day the uses of iron are increasing. The 
continual cheapening of both steel and wrought iron by 
improved methods has caused their use in thousands of 
ways where wood was formerly demanded. These applica- 
tions are too well known, however, to need mentioning. 

Compounds of Iron 

20. Ferrous and Ferric Compounds. — Like several other 
metals, iron forms two general classes of compounds, the 
ferrous and ferric. The former are very unstable, and 



306 



MODERN CHEMISTRY 



when exposed to the air gradually change to the ferric. 
The reaction in the presence of free acid may be indicated 
thus : — 

2 FeS0 4 + H 2 S0 4 + O(air) = Fe 2 (S0 4 ) 3 + H 2 0. 

If there is no free acid present, a part of the ferrous salt 
is converted into the ferrie and another part into an 
insoluble basic salt. 

Experiment 181. — To distinguish between ferrous and ferric salts. 
Pulverize a crystal of ferrous sulphate and dissolve in a few cubic 
centimeters of water; divide into three portions. To one portion add 
promptly a few drops of ammonium hydroxide, to the second a few 
drops of potassium sulphocyanide solution, to the third a few drops of 
potassium ferrocyanide solution. Notice the results in each case. 
Now dissolve a little ferric chloride or nitrate in water, divide into 
three parts, and repeat the same tests. Compare results and tabulate 
as below. 





NH 4 OH 


KSCy 


K 4 FeCy 6 


FeS0 4 








Fe 2 Cl, 









Experiment 182. — To show the instability of ferrous salts. Quickly 
dissolve in cold water a little powdered ferrous sulphate, and divide 
into two parts. To one add a few drops of ammonium hydroxide, and 
note the color of the precipitate. Allow both portions to stand for 
some time. How does the greenish precipitate change in color? Into 
w 7 hat is it apparently converted ? Has the other portion changed any 
in appearance? How? Test it with potassium sulphocyanide or 
ammonia to learn what kind of a salt it 'is now. What are your 
conclusions? 

21. This experiment will show the tendency of ferrous 
salts. What is thus accomplished slowly by the action of 



IRON, NICKEL, COBALT 307 

atmospheric oxygen at ordinary temperatures is effected 
rapidly by nitric acid at the boiling point. As already 
seen, this acid is a strong oxidizing agent, readily giving 
up a part of its oxygen when heated, thus : — 

2 HN0 3 + (heat) = H 2 + N 2 4 + O. 

This nascent oxygen rapidly attacks any oxidizable 
substance that may be present. With ferrous chloride 
in the presence of hydrochloric acid, the following reaction 
takes place : — 

2 FeCl 2 + 2 HC1 + O = Fe 2 Cl 6 + H 2 0. 

Experiment 183. — To show the effects of nitric acid upon a ferrous 
salt. Dissolve a little ferrous sulphate in water, add a few drops of 
sulphuric acid aud then some nitric acid, and heat to the boiling point 
for two or three minutes. Does the solution change any in color? 
Xow test a part of it in two or three ways to learn whether it has been 
converted into a ferric salt. What are your conclusions ? 

22. Ferric salts, on the other hand, may be reduced to 
the ferrous by treatment with hydrogen sulphide or nascent 
hydrogen. The reaction may be shown thus : — 

Fe 2 Cl 6 + H 2 = 2 FeCl + 2 HC1. 

Experiment 181. — To prove the above statement. Put into a test- 
tube about 5 cc. of a solution of ferric chloride or nitrate and drop into 
it a good-sized granule of zinc. Xow add a little strong sulphuric or 
hydrochloric acid to cause a rapid evolution of hydrogen. In from 5 
to 7 minutes the yellow color of the ferric solution should have disap- 
peared. Test with ammonia or potassium sulphocyanide. What are 
your conclusions in the matter ? 

23. How to distinguish Ferrous from Ferric Salts. — 
From the preceding work it will be learned that ferrous 
salts in solution are usually colorless or very pale green, 
while ferric salts are light brown. Potassium sulpho- 



308 MODERN CHEMISTRY 

cyanide serves as the most delicate method of detecting a 
ferric salt, because even exceedingly small quantities will 
show the characteristic wine-red color ; with ferrous salts, 
however, it shows no reaction, hence will not indicate their 
presence. Ammonia gives precipitates with both classes 
of salts, deep reddish brown with the ferric, and greenish 
with ferrous. Potassium ferrocyanide and ferricyanide 
may also be used to distinguish between the two. 

24. Sulphates of Iron. — Ferrous sulphate, FeSO±, 7 H 2 0, 
the only common ferrous salt, is formed when iron is dis- 
solved in sulphuric acid. It is commonly known as cop- 
peras or green vitriol, and occurs in light green crystals. 
It is somewhat efflorescent, and gradually gives up its 
water of crystallization, turning white and breaking up 
into a powder, anhydrous ferrous sulphate. It is used con- 
siderably in making black ink and dyes, also as a deodorizer 
and disinfectant. 

Experiment 185. — To show one method of making ink. Prepare a 
strong solution of ferrous sulphate, and add to it a little of another 
solution made by soaking some powdered nutgalls in water. Notice 
the bluish black color obtained. Allow it to stand a few minutes, and 
notice whether the color deepens. Now add to the solution a few 
drops of a solution of oxalic acid. What happens? This suggests a 
method for removing ink stains without injuring the fiber of the paper 
or cloth. 

25. Ferric Chloride, Fe 2 Cl 6 . — This is a brownish yel- 
low salt, which rapidly absorbs moisture when exposed to 
the air. It is obtained when iron is treated with aqua 
regia or dissolved in hydrochloric acid with the addition 
of a crystal of potassium chlorate. It has little use except 
in the laboratory. 

26. The Sulphides. — Ferrous, FeS ; Ferric, FeS 2 . The 
former is a dark gray substance somewhat resembling cast 



IRON, NICKEL, COBALT 309 

iron. It is made by fusing together, in the proportion of 
their atomic weights, iron and sulphur, and is used exten- 
sively in the laboratory for making hydrogen sulphide. 
Ferric Sulphide is the native ore, pyrite, or "fool's gold." 
It is of a brassy yellow color, and frequently occurs in 
beautiful cubes or modified forms of the cube. It is very 
abundant, but has little use except in the preparation of 
sulphur dioxide for the manufacture of sulphuric acid. 

27. The Oxides. — Ferric, Fe 2 O s . This is met with in 
the ore, hematite, already mentioned. It is also formed 
when iron is exposed to moisture, and is known as rust. 
In the hydrated form, Fe 2 (OH) 6 , ferric hydroxide, it is 
formed when any ferric solution is treated with ammonia. 
As a reddish brown precipitate it has already been seen in 
several of the preceding experiments. It is sometimes 
used as an antidote for arsenic poisoning. Magnetite, 
Fe 3 4 , is regarded as a mixture, FeO, Fe 2 3 , the first 
being the umrmport&nt protoxide. The greenish precipitate 
obtained in some of the preceding experiments by adding 
ammonia to a solution of ferrous sulphate is ferrous 
hydroxide, Fe(OH) 2 , or FeO, H 2 0; that is, the hydrated 
form of the protoxide, FeO. 

Nickel : Ni = 58.7 

28. Distribution. — Like iron, nickel is never found pure 
except in meteorites, of which, as already stated, it often 
constitutes from 5 to 7 per cent. Its ores are fairly well 
distributed, but are nowhere in great abundance, and with 
them are alwaj^s associated cobalt and iron. 

29. Characteristics of Nickel. — Nickel is a silvery white 
metal with the faintest yellow tinge ; it is susceptible of a 
very high polish and does not tarnish in the air. It is 



310 MODERN CHEMISTRY 

very hard, melts at about white heat, may be welded like 
iron, is magnetic, and becomes brittle like cast iron upon 
the addition of such impurities as cast iron always con- 
tains — carbon and silicon. Its density is but little greater 
than that of iron. It is soluble in nitric acid. In most 
respects, therefore, it is very similar to iron, and strikingly 
different in one respect only. 

30. Uses. — Nickel is used very extensively in alloys, 
among them being certain coins; in german silver, con- 
sisting of nickel, zinc, and copper ; and with steel for 
armor plating in making what is known as Harveyized 
steel, noted for its hardness and toughness. 

Nickel is also used largely in plating various articles of 
ornament and utility. 

31. Compounds. — The general color of the more com- 
mon nickel salts is green. Among these may be named 
the nitrate, Ni(N0 3 ) 2 , chloride, NiCl 2 , sulphate, NiS0 4 ; 
also Ni(OH) 2 , nickel hydroxide. This last may be pre- 
pared from a solution of any of the preceding salts by 
adding a few drops of ammonia or caustic potash. 

Experiment 186. — To 3 or 4 cc. of a solution of any of the above 
salts, add a little caustic soda. Describe the precipitate that forms. 
Test its solubility in hydrochloric acid. Write the two reactions 
taking place. 

32. A fifth compound which may be mentioned is the 
sulphide, NiS. It is prepared, as is the sulphide of other 
kindred metals, by adding ammonium sulphide to a neutral 
or alkaline solution of any nickel salt. 

Experiment 187. — Add a little ammonium sulphide to 4 or 5 cc. 
of a solution of any nickel salt. Describe the nickel sulphide that 
forms. Test its solubility in hydrochloric acid. Also in aqua regia. 
Write the reactions. 



IB OX* NICKEL, COBALT 311 

33. Nickel salts fused in a borax bead impart to it a 
smoky yellow or brown color according to the amount of 
the nickel present. 

Experiment 1SS. — Make a small loop in the end of a platinum 
wire, heat it in the burner flame and dip into some powdered borax. 
Now hold again in the flame until the borax which swells up at first 
has formed a clear transparent glassy bead. Dip into a solution of 
some nickel salt, or touch it to a tiny particle of nickel salt and fuse 
again. If you use the solution, it may be necessary to dip the bead 
several times. Note the color imparted. 

Experiment 189. — To find the composition of a coin. Put a 
••nickel" into an evaporating dish and treat with warm nitric acid for 
a few minutes. Remove the coin and add a few cubic centimeters of 
water. Pass a current of hydrogen sulphide through the solution for 
several minutes, and filter out the black precipitate. After washing 
it. punch a hole in the bottom of the filter and wash the precipitate 
through into a beaker with a little nitric acid. Heat until it dissolves. 
What colored solution is obtained? What metal is indicated by this 
color ? Add ammonia till alkaline ; is a deeper blue solution obtained 't 
What metal is it? 

Boil nearly to dryness the filtrate from the black precipitate above ; 
note the color. Does this indicate any salts with which you are 
familiar? Make a borax bead as directed in the preceding section 
and test the solution ; what are your conclusions ? 

Of what two metals is the coin -composed? If you can obtain one 
of the lighter-colored pennies seen occasionally, test it in the same 
way. 

Cobalt: Co = 59 

34. Characteristics. — This is a somewhat rare metal that 
is usually found associated with nickel. It is very similar 
to iron and nickel in its characteristics. being: steei-o:rav in 
color, very hard, magnetic, and of about the same melting 
point. It is permanent in the air. The metal itself has 
no application in any of the arts. 

35. Compounds. — Cobalt forms salts with the three 
common acids : the nitrate, Co(X0 3 ) 2 ; chloride. CoCl 2 *. 



312 MODERN CHEMISTRY 

and sulphate, CoS0 4 . These are all some shade of red 
in color, but when heated so as to lose their water of 
crystallization they become blue. 

36. The hydroxide and sulphide are prepared just as the 
similar compounds of nickel are. 

Experiment 190. — Prepare the last two as you did the correspond- 
ing compounds of nickel in Experiments 186 and 187, and test their 
solubility in the same way. 

37. None of the above has much use except occasionally 
in the laboratory. There are one or two others, however, 
which have extensive application in the arts. Among 
these may be named 

38. Smalt, a silicate of cobalt. When fused with glass 
or pottery ware this imparts a beautiful blue color, and is 
largely used for that purpose. It may be illustrated in the 
following experiment : — 

Experiment 191. — Prepare a borax bead as you did for nickel, 
and fuse with it some salt of cobalt. Note the color imparted. If it 
is too dark to recognize, it is because too much cobalt has been intro- 
duced. Break out the bead, and repeat the experiment, using less of 
the compound. 

39. Sympathetic Inks. — They are inks which under 
ordinary circumstances are invisible, or nearly so, on paper ; 
when heated or treated by some other method they become 
legible. Many of these have some compound of cobalt as 
their basis. 

Experiment 192. — Mix a solution of some compound of cobalt 
with one of ferrous sulphate. Using this as an ink, write with it upon 
paper, and when the inscription is dry heat it. Do you obtain a dis- 
tinct green color, though before it was nearly invisible ? In the same 
way try potassium iodide mixed with the cobalt solution. Results ? 



IRON, NICKEL, COBALT 313 



SUMMARY OF CHAPTER 

Iron, Nickel, Cobalt. 

Occurrence — Wherein are iron and nickel similar. 

History of some large meteorites. 

Some important ores of iron. 
Names and formulae. 
Localities where found. 
Reduction of iron ores — Description of blast furnace. 

Drawing of essential features. 

Method of charging the furnace. 

Chemical action that takes place. 

Plan of molding pig iron. 
Varieties of iron — Three. 

How different in composition and properties ? 

Description of the puddling furnace. 
Chemical action. 

Meaning of term bloom. 

Description of the converter. 

Explanation of the chemical changes. 

Plans used for phosphorus-bearing iron ores. 
Tempering steel. 

Meaning of the term. 

Process used. 
Characteristics — Compare iron and nickel as to 

Color. Susceptibility of polish. 

Hardness. Permanency in the air. 

Melting point. Several other similarities. 

Magnetic properties. 
Uses — Compare nickel and iron. 
Compounds — Classes of iron compounds. 

Compare them as to stability. 

Plans for distinguishing the two. 

Method of converting each into the other. Explain 
the chemical action in each case. 

Names of three or four compounds of iron and their 
uses. 

Compare the compounds of nickel and cobalt in color 
and method of preparation, 



CHAPTER XXVI 

THE PLATINUM GROUP 

Platinum : Pt = 195 

1. Where obtained. — Platinum is a rare metal, usually 
found uncombined, but almost always associated with 
iridium, and smaller quantities of palladium and osmium. 
The greater portion of the commercial supply comes from 
the Ural Mountains in Russia, though small quantities 
have been obtained in California, Arizona, and some parts 
of South America. 

2. Characteristics. — Platinum is a hard, silvery white 
metal, unaffected by the air at any temperature. It is 
somewhat malleable, but becomes less so if alloyed with a 
small per cent of iridium, though by this means its hard- 
ness is increased. It is a very dense metal, with a specific 
gravity of 21.5, osmium, the heaviest metal known, having 
a density of only 22.5. The melting point of platinum is 
about 1900° C, and it can be fused only by such intense 
heat as that of the oxyhydrogen blowpipe, or acetylene 
blast lamp. Like gold, it is soluble only in nitro-hydro- 
chloric acid, forming therewith platinic chloride, PtCl 4 . 

3. Property of occluding Gases. — The most remarkable 
property of platinum is that of occluding or absorbing 
various gases within its pores. It is estimated that at 
ordinary temperatures it will absorb 200 times its own 
volume of oxygen. In the spongy form, that is, when finely 
divided, as in the case of a metallic precipitate, the power of 

314 



THE PLATINUM GROUP 315 

absorption is especially striking. If a current of hydrogen 
be directed against the platinum sponge, so rapid will be 
the absorption that almost instantly the metal will become 
red hot, and in two or three seconds the jet will be ignited. 

Experiment 193. — Repeat Experiment 23 with hydrogen. 

4. If into a jar of hydrogen and oxygen, mixed in the 
proportion of two to one, a platinum sponge be introduced, 
the gases will be made to unite with explosive violence. 
This power of occlusion may be seen in the case of certain 
other gases, as ammonia and common coal gas. 

Experiment 194. — Support upon a ring-stand a small flask con- 
taining some strong aqua ammonia ; warm it gently so as to secure a 
constant and rapid evolution of gas from the liquid. Now heat to 
bright redness in the Bunsen flame a spiral of platinum wire, made 
by coiling it about a small glass rod, and hold it in the neck of the 
flask. The wire will continue to glow, and the intensity of the heat 
will often be increased. 

Take this same platinum coil and flatten it a little so as to bring the 
parts of the spiral closer together; hold it in the Bunsen flame until 
red hot, then turn off the gas. When the redness has just disappeared 
from the wire, again turn on the gas. The wire' will quickly grow red 
again, and in two or three seconds will re-ignite the escaping gas. 
This may be repeated over and over again. The same may be tried 
with a spirit lamp. 

5. Platinum Alloys. — Platinum readily alloys with 
lead, silver, antimony, and other metals which are easily 
reduced from their compounds ; hence it should never be 
strongly heated in contact with them. It is likewise 
injured by heating in a smoky flame, or by placing upon 
red-hot charcoal, which blisters the surface. Platinum 
vessels are usually cleaned by fusing in them for a few 
minutes some acid potassium sulphate, KHS0 4 , and are 
polished by rubbing gently with a little fine sea-sand. 



316 MODERN CHEMISTRY 

6. Uses. — The rarity of the metal and the long, com- 
plicated processes involved in preparing it in the pure 
form, make it almost as expensive as gold. It is worth 
from 50 cents to 75 cents per gram, or about $300 per 
pound. It is made into wire, foil, and various articles 
for use in the chemical laboratory, such as crucibles, 
dishes, tips of forceps, etc. To the chemist it is simply 
indispensable in analytical work. 

SUMMARY OF CHAPTER 

Names of the elements in this group. 
Source of the supply of platinum. 
Characteristics of platinum. 

Experiments that illustrate these. 
Alloys of platinum. 
Uses and value of the metal. 
Compare with metals studied previously as to 

Color. 

Melting point. 

Density. 

Tendency to oxidize. 

Power of occluding gases. 

Solubility in acids. 



CHAPTER XXVII 

CHROMIUM AND ITS COMPOUNDS 

Chromium : Cr = 52 

1. Where found. — Chromium is a rare metal which 
received its name from the Greek word, chromos, meaning 
color, and is so named because of the striking colors of 
most chromium compounds. It occurs chiefly in the 
Shetland Islands in the form of chromite, or chrome iron, 
Cr 2 3 , FeO, also written FeCr 2 4 . It is also found as 
crocosite, PbCr0 4 , in Siberia, Brazil, and the Philippine 
Islands. 

Compounds of Chromium 

2. Classes. — In the metallic form chromium has but 
little use. Its compounds, however, have various applica- 
tions. They may be divided into two important classes : 

3. Chromium as a True Metal. — Those in which chro- 
mium acts as a metallic element, with the power of 
replacing the hydrogen in acids to form salts. Of these, 
as in the case of iron, mercury, and others, there are two 
divisions, the chromous and chromic, but only the latter 
are important. As examples, we have chromic chloride, 
CrCl 3 , chromic nitrate, Cr(N0 3 ) 3 , and chromic sulphate, 
Cr 2 (S0 4 ) 3 . These as a rule are green in color, but the 
double sulphate of potassium and chromium, K 2 Cr 2 (S0 4 ) 4 , 
is violet. Solutions of the chromic salts are precipitated 
by caustic potash or ammonia, giving the hydroxide, 
Cr(OH) 3 . 

317 



318 MODERN. CHEMISTRY 

4. Chromium as an Acid Producer. — Those in which 
chromium serves as a non-metallic element, forming acids. 
Of these there are three classes, but only two merit notice, 
the chromates and the dichromates. The chromates are 
based on the theoretical chromic acid, H 2 Cr0 4 , wherein 
the chromium atom is that which distinguishes the acid, as 
does the sulphur in sulphuric acid, H 2 S0 4 . The general 
color of the chromates is yellow, though there are some 
exceptions. The best-known example is potassium chro- 
mate, K 2 Cr0 4 . 

ExrERiMENT 195. — To prepare some other chromates. To 3 or 4 cc. 
of a solution of potassium chromate in a test-tube add a few drops of 
lead nitrate or acetate in solution. Notice the color of the lead chro- 
mate formed. In the same way prepare some barium chromate by us- 
ing barium chloride with the potassium chromate. Compare its color 
with the preceding. Now prepare some silver chromate by using silver 
nitrate solution with the potassium chromate. Note its appearance. 

5. Potassium Dichromate, K 2 Cr 2 7 , orange-red in color, 
is the best-known example of the dichromates. 

Tabular view of the compounds : — 
I. Chromium as a true metal : — 

1. Chromous. 

2. Chromic — 

a. Chloride, CrCl 3 . 

b. Nitrate, Cr(N0 3 ) 3 . 

c. Sulphate, Cr 2 (S0 4 ) 3 . 

II. Chromium as an acid former : — 

1. Chromates — 

a. Potassium, K 2 Cr0 4 . 

b. Lead, PbCr0 4 . 

c. Barium, BaCr0 4 . 

2. Dichromates — 

a. Potassium, K 2 Cr 2 7 . 



CHROMIUM AND ITS COMPOUNDS 319 

6. Conversion of One Class of Compounds into Another. — 

Though the chromium salts are stable, they may easily be 
converted from one into another. By adding a little acid 
and passing a current of hydrogen sulphide through a 
solution of potassium chromate, the latter is changed into 
a salt of the first class (chromic). The change of color 
to green indicates that the reduction has taken place ; at 
the same time free sulphur is precipitated. Thus : — 

2 K 2 Cr0 4 + 3 H 2 S + 10 HC1 = 4 KC1 + 2 CrCl 3 
+ 8 H 2 + 3 S. 

Experiment 196. To prove the above. — Put into a test-tube a 
few cubic centimeters of a solution of potassium chromate, and add 
a little hydrochloric acid. Now pass through the solution a current 
of hydrogen sulphide. What change in color is noticed ? Is the 
sulphur precipitated? 

7. Sulphur dioxide has a like reducing effect upon a 
chromate solution. 

Experiment 197. — Put into a test-tube 4 or 5 cc. of a solution of 
sodium sulphite, Xa 2 S0 3 , and a little hydrochloric or sulphuric acid. 
Xotice that sulphur dioxide gas is being liberated. Now add a little 
potassium chromate or dichromate. How does the chromium solution 
change in color ? If sodium sulphite is not to be had, fill a bottle with 
sulphur dioxide gas, pour in the dichromate, and shake. 

Experiment 198. To show the reduction of the dichromates to the 
chromic salts. — Put into an evaporating dish 10 or 15 cc. of a solution 
of potassium dichromate, add some hydrochloric acid, and boil a few 
minutes. The addition of a little alcohol will hasten the action. 
Xotice the change in color. What compound of chromium is probably 
formed ? 

8. The above experiments prove that either the chro- 
mates or dichromates may be reduced to salts of the first 
class. The reaction that takes place in the latter is as 
follows : — 

K 2 Cr 2 7 + 14 HC1 = 2 KC1 + 2 CrCl 3 + 7 H 2 + 6 CI. 



320 MODERN CHEMISTRY 

Experiment 199. — To 2 or 3 cc. of potassium chromate solution 
in a test-tube add a few drops of hydrochloric or nitric acid. How 
does its color change ? What other salt of chromium in solution does 
it resemble ? In like manner treat 2 or 3 cc. of potassium dichromate 
solution with a few drops of caustic potash or any alkali. Notice the 
change in color ; what chemical change has taken place ? 

9. It will be seen by the above experiments that the 
chromates and dichromates may readily be converted, the 
one into the other. The reactions taking place are shown 
thus : — 

2 K 2 Cr0 4 + 2 HC1 = K 2 Cr 2 7 + 2 KC1 + H 2 0, 

and K 2 Cr 2 7 + 2 KOH = 2 K 2 Cr0 4 + H 2 0. 

10. The Oxides of Chromium. — Or 2 O s and CrO v chro- 
mium sesquioxide and trioxide. The former is basic in 
properties, the latter acid. The former is green, and is 
used in imparting a green color to glass and enamel ; the 
latter is a dark red crystalline solid. 

Experiment 200. — Make a borax bead and dip it into a solution 
of some chromium salt, then fuse in the burner flame. If a good color 
is not secured the first time, repeat the operation. 

11. Chromium trioxide may be prepared by adding 
strong sulphuric acid to a saturated solution of potas- 
sium dichromate. After standing for some time, beautiful 
red needle-like crystals separate from the liquid, thus: — 

K 2 Cr 2 7 + H 2 S0 4 = 2 Cr0 3 + K 2 S0 4 + H 2 0. 

These cannot be filtered out by ordinary methods, as the 
trioxide is a strong oxidizing agent and readily gives up a 
part of its oxygen to any organic compound, itself being 
changed into the sesquioxide, thus : — 

2Cr0 3 =Cr 2 3 + 3 0. 



CHROMIUM AND ITS COMPOUNDS 321 

12. Chromium trioxide is theoretically the anhydride 
of chromic acid, H 2 Cr0 4 , and seemingly ought to produce 
it when added to water, thus : — 

Cr0 3 + H 2 = H 2 Cr0 4 . 

But the action is merely one of solution, and the acid is 
not formed. 

13. Chromic Hydroxide, Cr(OH) 3 . — This is a green 
precipitate formed when ammonia or caustic potash is 
added to any chromic salt, as the chloride or sulphate. 

CrCl 3 + 3 KOH = Cr(OH) 8 + 3 KC1. 

14. Uses of the Compounds. — Some of the uses of chro- 
mium compounds, among others those of the sesquioxide 
and of lead chromate, have been mentioned. Both the 
chromate and di chromate of potassium are used as reagents 
in the laboratory, and in the arts for dyeing and calico 
printing. If the reaction, 

K 2 Cr 2 7 + 8 HC1 = 2 KC1 + 2 CrCl 3 + 4 H 2 + 3 0, 

is studied, it will be seen that potassium dichromate, treated 
with hydrochloric acid, is a strong oxidizing agent. Each 
molecule gives up three atoms of oxygen. If no other salt 
is present, this nascent oxygen unites with the hydrogen 
in six additional molecules of hydrochloric acid, thus : — 

6 HC1 + 30 = 3 H 2 + 6 CI. 

Combining the last two reactions, it will be seen that we 
have the one given on page 319, showing the reduction of 
potassium dichromate to chromic chloride. How r ever, if 
any oxidizable salt be present, as, for example, a ferrous 
compound, the nascent oxygen readily converts it from 



322 MODERN CHEMISTRY 

the ferrous to the ferric condition. This is shown in the 
following reaction : — 

2 FeCl 2 + 2 HC1 + O = Fe 2 Cl 6 + H 2 0. 

On account of this property, potassium dichromate is fre- 
quently used by chemists in estimating the amount of iron 
present in a solution. 

Experiment 201. — To illustrate this use and the oxidizing power 
of potassium dichromate. Dissolve a little ferrous sulphate in a few 
cubic centimeters of water and add some hydrochloric acid. Now 
add gradually drop by drop a solution of potassium dichromate. 
Notice how the solution changes to green. Test a portion of it with 
potassium sulphocyanide and learn whether the solution has been 
oxidized to the ferric condition. What are your conclusions? Study 
some of the foregoing reactions and see whether you can determine 
why the solution became green. 

SUMMARY OF CHAPTER 

Origin of the term chromium. Why applied to this metal. 
Classification of the chromium compounds. 

Names and formulae of the most important. 
Relation of the classes of compounds. 

Method of converting those of second class to first. 

Indication of the change. 

Method of converting chromates into dichromates, and vice 
versa. 
Compare the two oxides in 

Appearance. 

Properties. 
Commercial uses of certain compounds. 

Chromium sesquioxide. 

Chrome yellow. 
Laboratory uses. 

What uses as a reagent. 

How used as an oxidizing agent. 

Experiment to illustrate. 



CHAPTER XXVIII 

MANGANESE AND ITS COMPOUNDS 
Manganese : Mn = 55 

1. Where found. — This is a somewhat rare metal, often 
associated with iron ores. The most abundant natural 
compound is the dioxide, Mn0 2 , known as pyrolusite. In 
the metallic form, manganese has little use, but some of its 
compounds are valuable. 

Compounds of Manganese 

2. Classes. — These may be classified as follows : — 

3. As a Metal. — Those in which manganese acts as 
a metal, that is, having the power of replacing hydrogen 
in acids. These may be divided into 

a. Manganous, 

b. Manganic, 

of which only the former are important. The most com- 
mon of these are manganous chloride, MnCl 2 , and manganous 
sulphate, MnS0 4 , both crystalline salts, pink in color. 
From these may be prepared the hydroxide, Mn(OH) 2 , 
by adding ammonia to a solution of either salt; also the 
sulphide, MnS, by adding ammonium sulphide. 

Experiment 202. — Using a solution of either manganous chloride 
or sulphate, prepare the hydroxide and sulphide as indicated above 
and describe their appearance. Test their solubility in hydrochloric 
acid. State results. 

323 



324 MODERN CHEMISTRY 

4. Manganese Dioxide. — In this connection we shall 
notice the most important of the oxides, Mn0 2 , man- 
ganese dioxide. It is a black compound, and is used in pre • 
paring oxygen, bromine, chlorine, and iodine. Notice the 
similarity in method of the last three. 



CI 

Br 

I 



Mn0 2 + 2 NaCl + 2 H 2 S0 4 = Cl 2 + MnS0 4 + Na 2 S0 4 4- 2 H 2 
" +2NaBr+ " = Br 2 + " + " -f « 

" +2NaI + " =1 a + " + " + " 



5. As an Acid Former. — Compounds in which man- 
ganese serves as an acid-forming element. Of these, there 
are two classes, 

a. Manganates, 

b. Permanganates. 

The first of these is based upon a theoretical acid, man- 
ganic, H 2 Mn0 4 , and none is of special interest to us. The 
best-known example of the second is potassium perman- 
ganate, KMn0 4 . 

6. Potassium Permanganate. — This is a dark purple 
crystalline salt, soluble in water. It is used frequently in 
the laboratory as a reagent, in a technical way for the 
estimation of iron in iron ores, and for the testing and 
purification of cistern water. Like nitric acid and potas- 
sium dichromate (see pages 88, 321), it is a strong oxidizing 
agent. When treated with hydrochloric or sulphuric acid, 
it gives up oxygen, thus : — 

2KMn0 4 + 6HCl = 2KCl + 2MnCl 2 + 3H 2 + 5 0; 
2 KMn0 4 +3 H 2 S0 4 =K 2 S0 4 + 2 MnS0 4 + 3 H 2 + 5 O. 

The nascent oxygen thus obtained may be used in oxidiz- 
ing ferrous salts to the ferric condition, or in destroying 



MANGANESE AND ITS COMPOUNDS 325 

(oxidizing) the organic matter contained in a solution. 
In the case of the iron the reaction may be shown thus : — 

10 FeS0 4 + 8 H 2 S0 4 + 2 KMn0 4 

= K 2 S0 4 + 2 MnS0 4 + 5 Fe 2 (S0 4 ) 3 + 4 H 2 ; 

or, the live atoms of oxygen set free as shoAvn above de- 
compose five additional molecules of sulphuric acid, thus : — 

5 + 5 H 2 S0 4 = 5 H 2 + — (S0 4 ) 5 . 

Then the five (S0 4 ) groups or ions unite with the 10FeSO 4 , 
forming the ferric salt, 5 Fe 2 (S0 4 ) 3 . Sometimes it is 
written thus : — 

10 FeO + 50 = 5 Fe 2 3 , 

which expresses in a simple form the same change from a 
ferrous to a ferric condition. 

Experiment 203. — To show the oxidizing power of potassium per- 
manganate. To a fresh solution of ferrous sulphate add one or two 
cubic centimeters of sulphuric acid, and then slowly, drop by drop, 
potassium permanganate until the solution just begins to turn pink. 
Now test it with potassium sulphocyanide or ammonium hydroxide. 
Have you obtained a ferric salt? In this connection study the preced- 
ing reactions. 

In the same way test some cistern water that has an offensive odor. 
Before adding the permanganate heat the water nearly to boiling. 
Does it lose its odor by this treatment ? In the same way try some 
cistern water discolored with cedar shingles; is the color removed? 
Try also a strong solution of logwood; can you remove the dark color? 

What instances can you give in which nitric acid has served as an 
oxidizing agent? Potassium dichromate? 

7. The sulphuric acid is added simply to dissolve a dark- 
colored precipitate that would otherwise form and obscure 
the results. In purifying cisterns, of course the acid can- 
not be used, but the brown solid in a short time settles to 
the bottom and remains there. The amount of organic 



326 



MODERN CHEMISTRY 



matter in cistern water may be learned by measuring the 
amount of potassium permanganate added before the water 
begins to turn pink. Sometimes a manganese solution or 
salt is proved by the color it imparts to the borax bead. 

Experiment 204. — Prepare a bead as in the case of nickel or 
cobalt, and fuse with some salt of manganese. Notice the beautiful 
color imparted. 

SUMMARY OF COMPOUNDS 



A true metal in 
its chemism. 



1. Manganous \ 



2. 



a. Chloride, MnCl 2 . 

b. Sulphate, MnS0 4 . 

c. Hydroxide, 

Mn(OH) 2 . 
[ d. Sulphide, MnS. 



Manganic, Dioxide, Mn0 2 . 
Manganates, not important. 
Permanganates, Potassium, 
KMn0 4 . 

Compare the above compounds with those of chromium 
and note the few differences. 



Class II [ 1 
An acid-forming j 2 
element. [ 



SUMMARY OF CHAPTER 

Occurrence of manganese. 

How associated. Chief ore. 

Classification of its compounds. 

Compare with the compounds of chromium, showing wherein 
similar and wherein different. 
Uses of certain compounds. 
Manganese dioxide. 
Appearance. 
What laboratory uses. 
What commercial uses. 
Potassium permanganate. 
Appearance. 

Laboratory uses. Experiments to illustrate. 
Practical uses. Experiment to illustrate. 



APPENDIX A 



QUALITATIVE ANALYSIS 

It is not intended in the following pages to give any- 
thing like a complete system of qualitative analysis. Such 
would be impossible, keeping within the necessary bounds 
of a high-school text. As a matter of reference, however, 
and to meet the demand of any who may care to pursue 
to some extent this line of work, the following brief 
outline is offered. 

The student has noticed already that a reagent which 
will precipitate some metals from their solutions may have 
no effect upon various other metals. Taking advantage 
of this fact, we are able to divide the metals into groups, 
and then to separate the members of these groups one from 
another. Accordingly, depending upon the reagents used 
for precipitating the metals, five divisions are usually 
made as follows : — 

Precipitated as chlorides, 
Group I PbCl 2 , Hg 2 Cl 2 , AgCl, 

by using hydrochloric acid. 



Group II 



2. Mercury 
(ons salts) 

3. Silver 
Antimony 
Tin 

Arsenic 
Mercury 

(ic salts) 
Copper 
Bismuth 
Cadmium 



Precipitated as sulphides 
Sb 2 S 3 , SnS or SnS 2 , etc., 
with sulphureted hydrogen. 
The first three are soluble in 
yellow ammonium sulphide or 
sodium sulphide ; the others, 
not. 

327 



328 



MODERN CHEMISTRY 



Group III 



Group IV 



Group V 



Iron 

Aluminum 
Chromium 
Cobalt 
Nickel 
Manganese 
. Zinc 

Calcium 
Strontium 
Barium 
Magnesium 

Lithium 
Ammonium 
Sodium 
Potassium 



The first three are precipi- 
tated as hydroxides with am- 
monia, and constitute division 
one of this group. The last 
four are precipitated by am- 
monium sulphide as sulphides. 

Precipitated as carbonates, 
CaC0 3 , SrCOg, etc., 
with ammonium carbonate 
from an alkaline solution. 

Not precipitated by any com- 
mon reagents. Most of them 
usually tested by color impart- 
ed to flame, or the spectrum. 



The General Plan. — Suppose now we have a solution 
which may contain salts of any or all of the above metals. 
By adding hydrochloric acid, those of the first group would 
be precipitated and their chlorides separated by filtering. 
The filtrate would contain the remaining four groups. 
This would now be treated with hydrogen sulphide, 
whereby the second group metals may be precipitated and 
filtered out. In a similar way the separation of the third, 
fourth, and fifth groups would be effected. All that re- 
mains is to separate the metals of each individual group 
and prove their presence by means of some distinctive 
test. 

Ionic Theory. — A clear understanding of the processes 
underlying any qualitative analysis is rendered much 
easier by what is known as the Ionic theory. It has long 
been observed that certain elements or groups of elements 



APPENDIX A 329 

always give the same distinctive tests with certain re- 
agents. For example, a silver solution gives the same 
characteristic precipitate with any soluble chloride, 
whether it be hydrochloric acid, sodium chloride, or any 
other. 

Suppose in analyzing an unknown solution we have 
found four bases and four acid radicals : each base might 
have been combined with each of the acid groups, making 
in all sixteen possible cases. Were we compelled to test 
for each one of these possible compounds, analysis would 
be very tedious ; but, as already stated, each base affords 
the same test as if it existed alone. 

It seems, therefore, that when substances are dissolved, 
they become more or less dissociated. For example, 
hydrochloric acid becomes largely broken up into hydro- 
gen and chlorine atoms ; potassium chlorate into potas- 
sium, K and C10 3 , groups. As the solution becomes more 
dilute, this dissociation as a rule increases. 

Ions. — These dissociated atoms or groups of atoms are 
called ions, and the process itself, ionization. They are 
regarded as being charged with electricity, and are of two 
kinds, anions or negative ions, and cathions or positive 
ions. The metals, ammonium, and hydrogen are cathions; 
the acid radicals and elements, like N0 3 and CI, and the 
group HO, hydroxy 1, are anions. This is often called the 
theory of electrolytic dissociation, and concisely stated is 
that when compound substances are dissolved in water, 
they are to a greater or less extent broken up into their 
constituent anions and cathions. 

Application of the Theory. — In the brief space of this 
text it is impossible to make application of the theory to 
any extent. For this the student is referred to Ostwald's 
Analytical Chemistry, translated by McGowan. An illus- 



330 MODERN CHEMISTRY 

tration may, however, make the theory somewhat clearer. 
Suppose we have a solution containing lead nitrate, Pb 
(N0 3 ) 2 , silver nitrate, AgN0 3 , and mercurous nitrate, 
HgN0 3 . According to the ionic theory, the solution 
contains, not molecules of the three compounds mentioned, 
but largely individual ions of Pb, Ag, Hg, and (N0 3 ); 
hence, tests need be made only for these four. Now, 
when we add dilute hydrochloric acid, we introduce two 
other ions, H and CI. When those of Pb, Ag, and Hg 
meet with the CI ions, compounds form, which in the main 
are insoluble in water, hence are not dissociated or broken 
up into ions, and therefore fall as precipitates. The same 
is true in any other chemical reactions. 

Details of the Work. Group I. — Take about two-thirds 
of the unknown solution, " Solution A," and add to it a 
little hydrochloric acid ; if any of the first group metals 
are present, they will come down as a white precipitate. 
Filter out and save the clear filtrate for work with the 
remaining groups. We will label this " Solution B." To 
be sure that enough hydrochloric acid has been used, add 
a drop or two to this filtrate. If it becomes turbid more 
must be added, and the whole solution again passed 
through the filter paper. Now wash the precipitate on 
the paper two or three times with cold water, and throw 
out the wash water. Next punch a hole in the bottom 
of the paper, and by directing a stream of water from 
the wash bottle upon the precipitate wash it through 
into a beaker. Do not use too much water, however ; 
usually 50 to 75 cc. will be sufficient. If the precipitate 
is not easily loosened by the stream of water, remove 
it with a spatula or stirring rod, and add it to what has 
already been washed into the beaker. Next, heat this to 
the boiling point and after a minute or two filter quickly. 



APPENDIX A 



331 



If any precipitate remains upon the filter, wash once or 
twice with hot water. 

Tests for Lead and Mercury. — Lead chloride is very 
soluble in hot water, and if it was present it will now be 
found in the filtrate. Test a portion of it with potassium 
dichromate, K 2 Cr 2 7 ; another portion, with potassium 
iodide, KI, or sulphuric acicl. The first two give dis- 
tinctive yellow precipitates, the third, a heavy white one, 
somewhat soluble in water, but almost entirely insoluble 
in alcohol. Any precipitate left on the filter paper above 
will contain the mercurous and silver chlorides, if any 
were present. The latter of these is very soluble in am- 
monia ; so pour upon the filter paper a few cubic centi- 
meters of ammonium hydroxide. If mercury is present, 
the precipitate will turn black, and further proof is un- 
necessary. 

TABLE I 

Separation of Lead, Mercury, and Silver 



To the unknown solution, 
add HC1, filter out the chlo- 
rides, and wash the precipi- 
tates. Save the filtrate for de- 
termining metals of Group II 
and those following. Mark it 
" Solution B." Transfer the 
precipitates to a beaker ; add 
H 2 0, and boil. Filter, and 
wash with hot water, if any 
precipitate remains. Test 
filtrate for Pb as in 1. De- 
termine Hg and Ag in the 
precipitate as in 2 and 3. 



1. Test the hot water filtrate for 
Pb with K 2 Cr 2 7 , KI, and H 2 S0 4 . 
For results, see preceding work. 



2. To the precipitate left undis- 
solved by the hot water, add NH 4 OH. 
If it turns black, mercurous salts are 
indicated. Test filtrate that runs 
through, for Ag by 3, below. 



3. To the filtrate from 2, above, 
add HXOo till odor of NH 3 is no 
longer perceptible. A white precipi- 
tate indicates silver. 



332 MODERN CHEMISTRY 

Test for Silver. — To determine whether silver is 
present put the ammonia solution that has just filtered 
through into a test-tube and add nitric acid until no 
longer alkaline. This will be known by the absence of 
the odor of ammonia. If there is any silver present, a 
white precipitate will form, which may again be dissolved 
by adding ammonia. 

Group II. — Through " Solution B," the filtrate from 
the chlorides of the first group, pass a current of hydro- 
gen sulphide, until, after shaking the solution, the odor 
of the gas is very perceptible. Any metals of this group 
will now be in the form of sulphides. Warm somewhat 
to collect the precipitates, and filter quickly. Preserve 
the filtrate, " Solution C," for determining metals of the 
third and succeeding groups. 

Now wash the precipitates left on the filter and reject 
the w^ash water. Transfer the precipitates to an evapo- 
rating dish and add a few cubic centimeters of yellow 
ammonium sulphide or sodium sulphide in solution, and 
warm gently for several minutes. This will dissolve the 
sulphides of division 1 of this group, that is, those of 
arsenic, tin, and antimony ; while those in the second 
division, mercuric salts, copper, bismuth, cadmium, and, 
as lead chloride is somewhat soluble in water, sometimes 
lead, will remain as precipitates. It should be stated, 
however, that copper sulphide is partially soluble in 
strong yellow ammonium sulphide, hence, when its pres- 
ence is suspected from the color of the original solution, 
it is better to use sodium sulphide to separate division 
one from two. 

When the sulphides have been digested as stated, 
filter and wash the remaining precipitate with water to 
which a drop or two of ammonium sulphide has been 



APPENDIX A 333 

added. Save the filtrate to test for arsenic, tin, and 
antimony. 

Test for Mercury. — Transfer the precipitates of mer- 
cury, copper, etc., to a beaker, add a few cubic centi- 
meters of dilute nitric acid, and boil- All the sulphides 
will dissolve except that of mercury, which will remain as 
a heavy black residue. Disregard any dark-colored par- 
ticles that remain floating upon the liquid, for they 
consist merely of sulphur colored with small portions of 
the black sulphides not yet dissolved. The student can 
prove this by collecting them upon a small loop in a 
platinum wire and igniting in the bunsen flame. The 
mass will burn with characteristic flame and odor. The 
indications of mercury shown by the black residue may 
be verified by filtering out, washing, and dissolving in 
aqua regia. Boil to dryness, take up with water, and test 
one portion with stannous chloride. A white precipitate, 
turning gray when heated, or when more of the stannous 
solution is added, is distinctive. Test another portion 
with potassium iodide, adding a drop at a time. A bright 
red precipitate, soluble in excess of the reagent, should 
form. 

The filtrate from the mercuric sulphide, containing 
copper, bismuth, etc., should be boiled nearly to dryness, 
and water added to dissolve the salts. Before proceeding 
farther, it is always better, if lead has been found in the 
first group, to test a small portion of this solution in water 
in a test-tube with sulphuric acid and a little alcohol 
added. If a precipitate of lead sulphate forms, it must 
be removed in the same way from the whole solution, 
using very little sulphuric acid. 

Test for Copper. — Now add ammonia to the solution, 
from which you have removed the lead, until alkaline. 



334 MODERN CHEMISTRY 

If the solution turns darker blue, copper is indicated ; 
at the same time bismuth will come down as a fine white 
precipitate. As the quantity of bismuth in solution is 
usually small, the student must be careful not to overlook 
it ; at the same time he must not mistake for bismuth a 
fine sediment sometimes carelessly allowed to collect in 
the reagent bottle used for ammonia. 

Test for Cadmium. — To determine whether cadmium is 
present, after filtering out the bismuth, add to the blue 
solution potassium cyanide in solution, drop by drop, until 
the blue color has entirely disappeared ; then pass a 
current of hydrogen sulphide, by which the cadmium, if 
present, will be precipitated as a bright yellow sulphide. 

Tests for Arsenic, Tin, and Antimony. — For separating 
and determining the presence of arsenic, tin, and antimony, 
various plans have been suggested, but nearly all are 
more or less tedious and require considerable care. The 
following plan, perhaps, is as satisfactory as any. To the 
ammonium sulphide solution of these metals, saved above, 
add dilute hydrochloric acid till the solution is no longer 
alkaline. The three metals will again be precipitated as 
sulphides. If the precipitate is pale yellow, or nearly 
white, and small in quantity, it probably consists mainly 
of sulphur, and none of the metals need be sought. If it 
is dark colored, gold or platinum may be present, or if 
copper has been found in the other division of this group, 
and ammonium sulphide was used instead of sodium sul- 
phide, the precipitate may be only copper. Filter, and 
throw out the filtrate, as it contains no metals. Wash 
the precipitate, as usual, and transfer it to a beaker. 
Now add a little strong hydrochloric acid and warm 
gently ; the sulphides of antimony and tin will dissolve, 
but the arsenic will be unaffected. Filter, and test the 



APPENDIX A 335 

filtrate as follows : put into it a bright iron wire or 
nail, and after warming gently let it stand about fifteen 
minutes. The antimony is reduced to the metallic form, 
and the stannic chloride to the stannous. Filter or de- 
cant and test the solution for tin with mercuric chloride. 
The results are those given in testing for mercury 
with stannous chloride in the other division of this same 
group. 

Wash thoroughly the precipitated antimony, and add 
to it a little strong hydrochloric acid and a few drops of 
nitric acid. The antimony will dissolve. Boil the solution 
nearly dry and add water. A white precipitate indicates 
antimony, which may be verified by passing a current of 
hydrogen sulphide through the solution. An orange- 
colored precipitate will result. 

The arsenic left undissolved by the hydrochloric acid 
above may be tested in several ways. Transfer the ar- 
senic sulphide to a beaker, add to it some strong nitric 
acid, and heat. The arsenic will dissolve. Now fill a 
test-tube about half full of a solution of ammonium 
molybdate, add to it a few drops of the arsenic solution 
prepared above, and boil. A yellow crystalline precipi- 
tate indicates arsenic. 

Sometimes the following method works satisfactorily. 
After adding concentrated hydrochloric acid to dissolve 
the precipitates of antimony and tin sulphide obtained 
from the ammonium sulphide solution, decant the clear 
solution into a test-tube. Now slowly pour hydrogen 
sulphide water down the inside of the tube. Presently 
the antimony will begin to precipitate, forming an orange- 
colored ring of the sulphide. Continue adding the hy- 
drogen sulphide solution, when above the antimony a 
ring of yellow stannic sulphide will form. 



336 



MODERN CHEMISTRY 
TABLE FOR GROUP II 



>» tc o3 



H o 



t* o 



•^ fc ^ 

c$ _, ^ 

H-.S fe 



"si 

CO O 



w o 



p< 



* 8 

CD 5h 

5- d 

O C3 



— < rj -^ CM 

CO , 

o CD ^ 

a § .S>«8 

' 2 ^ "§ 

o g is 

? z *** 

5 ^ cS 

-3 s 



o 



ffl 



O co 



c8 -d 

S-i CO 

O 03 



2 
O 

*3 CO 
rd CD 

is 



CD CD 
c8 



> -4_3 ^) 



o c3 



CD CD 

cd ^ 

^ CD 

Q c3 

E o 

'$ £ 

c3 



CD 

■+■> S-i 

53 £- 

CD CD 

-d ^ 

^ CO 

*o ^ 



n3 cd 

§ 3 

CO -4J3 
CD *o 

Pm &• 

CD 

. rd 

s ^ 

s ^ 

c5 d 

c3 






bp 5 

§ 2 



co q=i 

c3 






a. Put into the filtrate a bright iron 
wire or nail and let stand about 15 minutes. 
A black scaly precipitate of antimony forms. 
Filter out and test by b. To the nitrate 
add HgCl 2 , drop by drop, as a test for the 
tin. 



b. Dissolve the precipitated antimony in 
aqua regia, boil nearly dry, and add water. 
A white precipitate indicates antimony, 
verified by H 2 S, which gives orange-colored 
precipitate. 



c. Heat the precipitate of arsenic sul- 
phide with a little nitric acid and add some 
ammonium molybdate solution. A yellow 
crystalline precipitate will indicate arsenic. 





a 



£2 
.5 a 

-S 3 * 
s £ - 

a ^ 

CD 



^ a. 

5 ® 

5-i C3 

CD "*J 



^5 33 
Ph o 

s *° 



T3 






ffi CD 



- :3 cd 

O 



B, ^< s 



CD CD 



cq 



B S 



.« S3 -d ^ .05 htm' 

5 - s fl 



.T3 O 



A' 



HH «4-l 
MH t— I 



«. Deep blue color indicates 
copper. 



b. White precipitate indicates 
bismuth. To verify, filter out, 
dissolve in II CI, boil nearly dry, 
and add H 2 0. While precipitate 
forms, filter. 



c. After filtering out the bis- 
muth, add KCy solution till the 
blue color has disappeared. Pass 
a current of H 2 S. A yellow pre- 
cipitate indicates cadmium. 



Group III. — Like Group II, this is also usually separated 
into two divisions for convenience in analysis. The first 
includes iron, aluminum, and chromium, precipitated by 
ammonia ; the second, manganese, zinc, nickel, and cobalt, 
with ammonium sulphide as the precipitant. 



APPENDIX A 337 

To " Solution C," the filtrate saved from Group II, after 
filtering out the sulphides, add a few drops of nitric acid 
and boil a short time. Now add ammonium chloride, 
NH 4 C1, and ammonium hydroxide till alkaline. The 
latter reagent precipitates the metals of the first division 
as hydroxides. Warm the solution, filter and wash as 
usual. Save the filtrate for the second division of this 
group and the succeeding groups. Transfer the hydrox- 
ides of iron, chromium, and aluminum to a beaker, add 20 
or 25 cc. of strong potassium hydroxide solution, and boil 
several minutes. This will dissolve the aluminum and 
leave the others unchanged. Filter and wash. To a por- 
tion of the filtrate, after acidulating with hydrochloric 
acid, add ammonia till alkaline. A white, flaky, some- 
times starchy precipitate indicates aluminum. 

Test for Iron. — Take a portion of the iron and chro- 
mium precipitate left undissolved and add hydrochloric 
acid. Test the solution obtained for iron, by using either 
potassium sulphocyanide, KSCy, or potassium ferro- 
cyanide. 

Test for Chromium. — Next, take a rectangular piece of 
platinum foil and bend up the sides so as to form a small 
boat or pan. A piece of broken porcelain dish may serve 
the same purpose, but more heat will be needed. Put into 
the boat the remaining iron and chromium precipitate, add 
an equal amount of potassium nitrate, KN0 3 , and as much 
sodium carbonate, Na 2 C0 3 , and heat red hot until the 
whole mass has fused well together. Upon cooling, if 
chromium is present, it will assume a yellowish appearance. 
Put the boat and contents into a beaker containing a little 
water and dissolve the mass. Acidulate the solution with 
acetic acid and test a portion with silver nitrate. A brick 
or blood red precipitate of silver chromate, Ag 2 Cr0 4 , indi- 



338 MODERN CHEMISTRY 

cates the presence of chromium. Test another portion 
with lead acetate, Pb(C 2 H 3 2 ) 2 . 

Tests for Nickel and Cobalt. — To the filtrate saved for 
the second division of this group, add some ammonium 
sulphide. If precipitates of light color are obtained, nickel 
and cobalt are not present, as their sulphides are black. 
If nickel is present, the filtrate will often be of a dark- 
brown color, which is apt to lead the student to think the 
solution is not filtering well. Disregard this, mark it 
" Solution D," and save for work with the fourth group. 
After washing the precipitates, transfer them to a beaker 
and treat with dilute hydrochloric acid ; the sulphides of 
zinc and manganese will dissolve, while those of nickel and 
cobalt will remain as a black residue. Filter and wash. 
Test the black residue with the borax bead ; cobalt gives 
the well-known blue in the oxidizing flame, and nickel, 
yellow to brown or black, according to the amount intro- 
duced into the bead. If both metals are present, the cobalt 
blue will obscure the brown, and further tests are neces- 
sary ; for these the student is referred to any manual on 
qualitative analysis. 

Tests for Zinc and Manganese. — To the solution sup- 
posed to contain zinc and manganese, after boiling for two 
or three minutes, add caustic potash till strongly alkaline. 
Allow it to stand for some time, for manganese precipitates 
slowly. If present, it may be filtered out and the precipi- 
tate tested with the borax bead. It imparts a beautiful 
amethyst color. Acidulate the filtrate with acetic acid and 
add ammonium sulphide till alkaline. A white precipitate 
indicates zinc. This is usually verified by heating on 
charcoal, moistened with a solution of cobaltous nitrate. 
A green mass is obtained; aluminum compounds treated 
in the same way give a blue mass. 



APPENDIX A 



339 



TABLE FOR GROUP III 




a. Test for Co with borax bead — 
blue. 



b. If Ni is present with no cobalt, borax 
bead will become yellow to brown in oxi- 
dizing flame. 

c. Add considerable excess of KOH 
and let stand for some time. A slow- 
forming precipitate indicates Mn. Filter 
and test filtrate by d. Verify Mn with 
borax bead. Amethyst color. 



d. Acidulate nitrate from c with acetic 
acid, add (NH 4 ) 2 S till just alkaline. 
Zinc forms white precipitate of ZnS. 



a. Filtrate may contain aluminum. Acidulate 
with HC1, add NH 4 OH or (NH 4 ) 2 C0 3 . A white 
precipitate of aluminum hydroxide shows the 
presence of the metal. 

b. Dissolve a small portion of the precipitate 
in HC1 and test for iron with KSCy, K 4 FeCy 6 , 
or NH 4 OH. 

c. Fuse the remainder in platinum dish with 
KNOo and Na 2 C0 3 . Dissolve in H 2 0, acidulate 
with HC 2 H 3 2 . Test for Cr with AgN0 3 , also 
with Pb(C 2 H 3 2 ) 2 . 



Group IV. — For this use " Solution D " saved from 
Group III. It is better to make a preliminary test be- 
fore proceeding with the whole. To do this, add to a 
small portion of the solution to be analyzed a little 
disodium phosphate. If a white precipitate forms, some 
of the metals at least are present, and all mast be tested 



340 MODERN CHEMISTRY 

for. If so, add to the whole ammonium chloride, am- 
monium hydroxide, and ammonium carbonate. A white 
precipitate may contain calcium, strontium, and barium in 
the form of carbonates ; filter and wash. Save the filtrate 
to test for magnesium and fifth-group metals. 

Test for Barium, Strontium, etc. — Transfer the precipi- 
tates to a beaker and dissolve in acetic acid. Test a small 
portion of this with potassium dichromate ; a light yel- 
low precipitate indicates barium, which may be verified by 
the flame test. If present remove it from the entire solu- 
tion by adding the dichromate and filtering. To the 
filtrate add caustic potash till alkaline and a little more 
potassium dichromate, when the strontium, if present, will 
be precipitated, as strontium chromate is insoluble in 
alkaline solutions. Remove this by filtering, and test 
the filtrate for calcium by adding ammonium oxalate. 
This gives a fine white precipitate of calcium oxalate. 
It is customary to verify the strontium by the flame 
test, as its salts impart a crimson color which is very 
persistent. 

There are other plans for effecting a separation of the 
metals of this group, of which the following is frequently 
used. After removing the barium, to a small portion of 
the filtrate add a little strong solution of calcium sul- 
phate. If a white precipitate forms, strontium is present 
and must be removed. Add to the remainder of the solu- 
tion a very little sulphuric acid ; the strontium will slowly 
precipitate. After a few minutes filter and test filtrate 
for calcium. To do this add sufficient ammonia to neu- 
tralize any excess of sulphuric acid present, and then add 
ammonium oxalate solution as in other methods. 

To a small portion of the filtrate saved above, after pre- 
cipitating the barium, strontium, and calcium with am- 



APPENDIX A 



341 



monium carbonate, add a little disodium phosphate ; a 
white precipitate which may form slowly will indicate 
magnesium. 

TABLE FOR GROUP IV 



ffl 


o 


DO 


+3 




o 


03 

o 


+3 


CD 
O 




« 


5 


PI 


od 




'A 


CD 


08 


.3 




3 


Eh 


•+3 


CD 

> 






(1) 


^ 


o 




K 




+j 




X 


-»-) 

fi 


CD' 


*8 




s 


£ 


> 
OS 

EC 


nri 




c8 


o 
XI 


08 






C/J 








HH 






u 




HI 


c/j 


X! 


CD 




p 

5 


XI 

+3 

CD 


e8 

a 

03 


08 

CD 

X> 

o 

■+3 

an 




g 


08 




CD 

-fc3 




o 




-4J 


art 




f_l 










<-M 


a 


^H 


& 




q 


® 




CD 
CD 


T— | 




Ph 


o 


2* 


o 


c5 


S-. 


~J2 

08 








1) 


s 


3D 


be 


an 


£ 


o 


^5 


w 




O 


£ 




c3 


■+3 

1 


£ 






H 




„ 


o 


o 

DO 




u 
en 


CL 

■+3 




^ 


+3 


1^ 


K 

fc 


08 

+^ 

CD 


X 
bfi 


■CD 

in 


o 


^3 


CD 

a 


S 


T3 




u 


n 


o 


a 




c8 


x 


e*H 


08 



1. To a small portion of the solution, add a little 
K 2 Cr 2 7 solution. If Ba is present, indicated by the 
forming of a light yellow precipitate, treat the whole of 
the solution in the same way, and filter. The Ba pre- 
cipitate may be verified by flame test. Test the filtrate 
by 2. 



2. Render the filtrate alkaline by adding KOH, and 
then add a little more K 2 Cr 2 7 . If strontium is present, 
it will be precipitated and may be filtered out. Or, from 
the filtrate from 1 above, the strontium may be removed 
by adding a little sulphuric acid. Let it stand a few T 
minutes and then filter, and test filtrate for Ca by 3. 
The precipitate may be. verified by flame test. 

3. When the barium and strontium have been re- 
moved, if the filtrate is not already alkaline, render it 
so by adding NH 4 OH. Then add ammonium oxalate ; 
white precipitate is distinctive of calcium. May be 
verified by flame test, orange-yellow. 

4. To a small portion of the filtrate saved from 
"Solution D," add a little disodium phosphate. A 
white precipitate indicates magnesium, distinctive in 
the absence of other metals of this group. Filter, and 
save filtrate for Group V, " Solution E." 



Group V. Sodium, Potassium, Lithium. — The salts of 
the metals of this group are all soluble in water, hence none 
of the reagents used in the previous steps of analysis pre- 
cipitate them. The flame, especially with the spectroscope, 
is usually all that is necessary for their identification. 



342 MODERN CHEMISTRY 

As sodium is so widely distributed, a slight test for it may 
nearly always be obtained, but the student must learn to 
disregard any except a decidedly strong indication. As 
already seen, if sodium is present, the potassium flame can 
be perceived only by making the observation through a 
sheet of cobalt glass. Before making these flame tests, 
boil the solution, saved from Group IV, to dryness, and 
heat gently until ammonia fumes are no longer driven off. 
Dissolve the residue in water, and acidulate with hydro- 
chloric acid. 

Sodium gives bright yellow flame, 
Potassium gives violet flame, 
Lithium gives bright red flame, lasting but a 
moment. 

Potassium may also be tested in another way. To the 
solution used in making the flame tests, add some platinic 
chloride in solution and a little alcohol. Allow it to stand 
for some time, stirring occasionally with a glass rod. A 
small quantity of a yellow precipitate of potassic-platinic 
chloride, K 2 PtCl 6 , is slowly deposited. A large watch 
crystal serves well for making this test. 

Test for Ammonia. — Ammonia must be looked for in 
the original solution, as so many ammonium compounds 
are used as reagents in making the analysis. Put a few 
cubic centimeters of the original solution into a beaker 
and add caustic soda or potash until strongly alkaline. 
Moisten the under side of a watch crystal with a drop of 
water and upon it place a short strip of red litmus paper. 
Put the crystal over the beaker, and warm the solution 
gently. If ammonia is present, it will be liberated by the 
non-volatile alkali added, and will turn the litmus paper 
blue, 



APPENDIX A 



343 



TABLE FOR GROUP V 



>> 
n 


j: 

3 


'5 


c5 
03 

E 


£5 


—^ 


u 


03 


cd 




> 


5 


3 


33 






o 


a 








"3 


3 


0O 

c8 

O 


E 



- 


__ 


p 




^3 


03 
Oh 

X 


o 


g3 


2 

s 


03 

-*-2 


0> 

n 


^) 


3 






+3 

bo 


> 


4h 

03 

H 






pq 













m 


5 


1— 1 














s 


Td 


i— H 


• 3 






1. Make test with platinum wire; Na gives 
yellow ; K, violet when alone ; Li, red. Use cobalt 
glass, if Na is present, to distinguish the violet rays. 

2. Sometimes K must be detected otherwise 
than by the flame test. To the acidulated solution, 
add PtCl 4 ; a yellow precipitate, slowly forming, 
indicates K. 

3. To a portion of the original solution, add 
caustic soda or potash till alkaline, and warm gently. 
Suspend a strip of red litmus in the fumes arising. 
If it turns blue, NH 4 is indicated. 



The five tables given above simply show in condensed 
form the methods already described ; and when the student 
has once seen the details and understands them, he will 
find the tables very convenient for rapid work. For a 
successful analysis, neatness is absolutely essential, and 
great care must be used in washing the precipitates so as 
to remove all of the metals contained in the filtrate. 

Detection of Acids. — As a rule, the beginner will only 
meet with a few of the more common acids, and these only 
will be noticed here. They may be placed in groups, 
somewhat as the metals are, according as they are affected 
by certain reagents. 

Group I. — This includes those which form a precipitate 
with barium chloride. The only one with which the 
student will meet often is sulphuric acid. As already 
seen, this gives with barium chloride, barium sulphate, 
BaS0 4 , insoluble in all acids. 

If the solution be neutral, phosphoric acid or the phos- 



344 MODERN CHEMISTRY 

phates also give a white precipitate with barium chloride ; 
but this is soluble in hydrochloric acid. After being thus 
dissolved, if the solution be made alkaline with ammonia, 
the precipitate will again fall. 

Sulphurous and thiosulphuric acids are usually put in 
this group. They may be easily distinguished, how- 
ever. To the solution add a little strong hydrochloric 
acid ; both sulphurous and thiosulphuric acid and their 
salts will give off fumes of sulphur dioxide which may be 
readily detected. The latter, however, at the same time, 
throws dow T n a milky or pale yellow precipitate of sulphur, 
while the former remains clear. The reaction is shown 
below : — 

Na 2 S0 3 + 2 HC1 = 2 NaCl + H 2 + S0 2 (sulphurous), 

Na 2 S 2 3 + 2 HC1 = 2 NaCl + H 2 + S0 2 + S (thiosulphuric). 

Group II. — This includes such as form no precipitate 
with barium chloride, but do with silver nitrate. The 
most common are : — 

Hydrochloric, HC1, curdy white precipitate, very solu- 
ble in ammonia. 

Hydrobromic, HBr, pale yellowish white precipitate, 
slowly soluble in ammonia. 

Hydriodic, HI, pale yellow precipitate, very slightly 
soluble in ammonia. 

Methods for testing each of these and its compounds have 
been given in the text, and the student is referred to them. 
Hydrogen sulphide, H 2 S, if free, is known by the odor. 
In the form of compounds, it may usually be detected by 
adding some acid and heating, whereby the gas is liberated 
and its characteristic odor becomes perceptible. 



APPENDIX A 345 

Group III. — Here belong those acids which form no 
precipitate with either barium chloride or silver nitrate. 
The only common one is nitric, but the salts of nitrous 
and chloric acids, HN0 2 and HC10 3 , especially those of the 
latter, have occasional use in the laboratory. A plan for 
testing and distinguishing between nitrous and nitric acids 
was given in the text. The following plan, however, 
usually works satisfactorily, and by some is preferred to 
the other. Into a test-tube, containing the solution to be 
tested, drop a crystal of ferrous sulphate, and then pour 
down the sides of the tube a few drops of strong sulphuric 
acid. A brown ring will form about the crystal of ferrous 
sulphate. 

The chlorates, for example, potassium chlorate, KC10 3 , 
heated with sulphuric acid, yield chlorine, and chlorine 
peroxide, a very explosive gas. Usually, if sulphuric acid 
is added to a crystal of the chlorate, a sharp explosion 
occurs, throwing the materials out of the tube. 

Group IV. — We might place here certain organic acids, 
which require special tests for identification. The only 
common one is acetic, HC 2 H 3 2 , though the student oc- 
casionally meets with one or two others. Acetic acid and 
its salts are tested by adding a solution of ferric chloride 
and boiling. The solution becomes a deep red color which 
may be destroyed by using hydrochloric acid or mercuric 
chloride. 

Oxalic acid, H 2 C 2 4 , might be placed here, though more 
properly in Group I, as its salts form a white precipitate 
with barium chloride in neutral or alkaline solutions ; 
this precipitate is soluble in hydrochloric acid, but not 
in acetic. 

Preliminary Work. — Before testing any solution for 
acids, the metals present should be determined, other- 



346 MODERN CHEMISTRY 

wise the student may be greatly misled. If any are 
present which would interfere with necessary tests, that 
is, if there are any which would form precipitates with 
the reagents necessarily used in making the acid tests, 
they must be removed before proceeding with the deter- 
mination. 

Again, if the unknown substance is in solution, it would 
be useless to look for the acids whose salts are insoluble 
in water. For example, if we have found lead or barium 
present in a given solution, obviously it would be unnec- 
essary to look for sulphuric acid. Hence a knowledge 
of the solubility of salts is very important, and the fol- 
lowing incomplete table is given, showing the solubility 
of a few of the more common salts: — 

Acetates, soluble in water. 

Bromides, nearly all soluble ; exceptions, those of first 
group metals and mercuric. 

Carbonates, only those of Group V, the alkali metals. 

Chlorides, nearly all, Group I excepted. 

Iodides, nearly all, Group I excepted, also certain 
iodides of bismuth and copper. 

Nitrates, all soluble. 

Nitrites, nearly all soluble. 

Phosphates, only those of Group V. 

Sulphates, many insoluble, such as those of barium, 
mercury, lead, and silver ; and calcium and stron- 
tium, nearly so. 

Sulphites, only those of Group V. 

Sulphides, only those of Groups IV and V. 

If the substance, of which the acid radical is to be 
determined, is not in solution, it is often of great advan- 
tage to make certain preliminary tests. Put a small por- 



APPENDIX A 3J:7 

tion of it into a test-tube and add a little strong sulphuric 
acid. Warm gently, and notice the color and odor of the 
gas obtained. The more common are shown below : — 

Acetates, odor of vinegar, no color. 

Bromides, sickening odor, brown color, resembling 
that of nitrogen tetroxide. Odor is more offensive 
and peculiarly irritating to the eyes. 

Carbonates, strong effervescence, no special odor, 
colorless gas. 

Chlorides, very irritating gas (HC1), colorless. 

Iodides, peculiar odor, resembling weak chlorine, 
violet color. 

Nitrates, very irritating gas, no color. 

Nitrites, irritating gas, brown in color ; not so offen- 
sive as bromine. 

Phosphates, no special action. 

Sulphates, no special action. 

Sulphites, suffocating gas (S0 2 ), colorless. 

Sulphides, offensive odor (H 2 S), colorless. 

Thiosulphates, suffocating gas (S0 2 ), colorless. 

The student must remember that these are merely pre- 
liminary steps and must be verified by distinctive tests 
already described. 



APPENDIX B 



SOME ADDITIONAL QUANTITATIVE WORK 

It is believed that all the quantitative work that the 
ordinary class can do has been introduced into the text. 
There may be occasions, however, when it will seem desir- 
able to vary the work or even to furnish more to certain 
students ; to meet such a demand, the following experi- 
ments are offered. 

1. To estimate Amount of Carbon Dioxide in any car- 
bonate soluble in acids. (Adapted from Fresenius.) 
Fit two small bottles with rubber stoppers and glass 

tubing, as shown in the 
figure. E is a short piece 
of rubber tubing which 
may be closed air-tight by 
means of a screw clamp. 
The carbonate to be used, 
calcite for example, 
CaC0 3 , is accurately 
weighed, placed in M, and 
covered with water. N 
is filled over half full of pure concentrated sulphuric acid. 
Find the weight of the whole, which should not exceed 
60 to 70 g., tighten the clamp at E, and test the appa- 
ratus to see that it is air-tight. 

Now by suction at (7, partially exhaust the air in N; 
this will have a like effect upon M, and upon readmitting 
the air to N the acid will be forced over into if. The 

348 




Fig. 05. 



APPENDIX B 349 

carbonate will thus be decomposed, and the carbon dioxide 
will escape into iV, being dried as it bubbles through the 
acid. When the carbonate has all been decomposed, and 
the evolution of gas has ceased, open the clamp at E, and 
by means of an aspirator or by suction remove the carbon 
dioxide from M and 2V", and when the apparatus has become 
cool, weigh again. The loss represents the amount of 
carbon dioxide expelled by the acid. 

Carbonate used (for example) . . 1.0 g. 

Apparatus and contents, say . .68.0 g. 

After decomposition by acid : — 

Total weight .... x g. 

Loss ...... 68.0 — x. 

C0 2 = 68 -xg. 

2. To determine the Water of Crystallization in a Com- 
pound. — This is usually done by heating a known weight 
of the compound, and noting the loss. To illustrate, put 
into a small porcelain crucible, the weight of which is 
known, about a gram of magnesium sulphate, and weigh 
carefully. Support the crucible in a clay triangle upon 
an iron ring-stand. With the Bunsen burner heat cau- 
tiously at first, increasing to red heat, cool slowly and 
weigh. Heat a second time four or five minutes and 
weigh again. Do this until two successive weighings 
show the same results, then calculate the per cent of 
water of crystallization. 

Tabulate results thus : — 

Weight of crucible + MgS0 4 . . . a 

Weight of crucible alone b 

Weight of MgS0 4 .... a-b 



350 MODERN CHEMISTRY 

After the second and third heating, when weight was 
the same : — 

Crucible + MgS0 4 c 

Loss . . . . . . . a — c 

In the same way try some other salt containing water 
of crystallization, as, for example, common alum or copper 
sulphate. 

3. Volumetric Composition of the Air. — As the air, dis- 
regarding the impurities and small portions of other gases 
present, consists of oxygen and nitrogen, we can remove 
the former by exploding with hydrogen 
and then measure the residue. For ex- 
ample, suppose we pass into the eudio- 
- air meter 20 cc. of air and then 10 cc. of 

hydrogen. We now have a total amount 
gases o £ gQ cc ^ . p ass an electric spark to ex- 
plode the hydrogen and oxygen. As 
two parts of hydrogen unite with one 
of oxygen, one-third of the loss would 
represent the oxygen, and the other two-thirds the hydro- 
gen, which has combined to form water. The residue 
will contain the nitrogen of the air and any excess of 
hydrogen. Take the quantities used above : — 




Air ..... 


20 cc. 


Air + H . 


30 cc. 


Residue after exploding 


18 cc. 


Loss .... 


12 cc. 



|- of loss = 4 cc, the oxygen of air used. 
20 cc. air = 4 cc. oxygen + 16 cc, nitrogen 
of air. 



APPENDIX B 351 

Let the student arrange his own apparatus for the above 
experiment, making all corrections necessary for accurate 
results, and prove the usual statement that air is one-fifth 
oxygen and four-fifths nitrogen. 

4. The Volumetric Composition of Ammonia. — The 
composition of ammonia may be determined, but the ex- 
periment requires time and is somewhat tedious. The 
plan is as follows : into a eudiometer, over mercury, 
introduce a few cubic centimeters of dry ammonia gas, 
and pass sparks from an induction coil for twenty or 
thirty minutes or until the volume of the ammonia seems 
no longer to increase. This, in accordance with the law 
of Gaj'-Lussac, should now be double what was intro- 
duced into the eudiometer. Next add sufficient oxygen 
to explode with the hydrogen obtained from the ammonia, 
and pass a spark. It is obvious, from the proportions in 
which hydrogen and oxygen combine, that two-thirds of 
the loss represents the hydrogen, which, subtracted from 
the volume of the gases after electrolysis, gives the 
amount of nitrogen contained in the ammonia. Take the 
following example : — 

Ammonia gas introduced 

Vol. of mixed gases after passing sparks, 
Oxygen added ..... 

Total amount ..... 

Residue after exploding 

Loss ...... 

Two-thirds of loss = hydrogen, which 
was obtained from the ammonia gas . 

Volume of mixed gases .... 
Subtract volume of H 

Volume of N . . . . . 4 cc. 



8 


cc. 


16 


cc. 


8 


cc. 


24 


cc. 


6 


cc. 


18 


cc. 


12 


cc. 


16 


cc. 


12 


cc. 



352 . MODERN CHEMISTRY 

This proves that hydrogen and nitrogen unite in the 
proportion of three to one to form ammonia ; furthermore 
we have seen that the four volumes of the mixed gases 
upon combining are condensed to two. 

Let the student arrange his own apparatus, taking 
every precaution to insure accuracy, and, using different 
quantities from those mentioned above, prove the truth 
of the preceding statements. 

5. Composition by Volume of Hydrochloric Acid. — This 
may be learned by the interaction of sodium and hydro- 
chloric acid, by which is formed common salt and free 
hydrogen. In order to lessen the rapidity of the reaction, 
an amalgam of sodium should be used. This may be 
prepared by putting a small quantity of mercury into a 
mortar, and then, by means of forceps, thrusting small 
pieces of sodium, one at a time, into the mercury. Do 
this until a pasty mass is obtained, which upon cooling 
becomes solid. In preparing the amalgam do not hold 
the face too close to the mortar, as the combination some- 
times takes place with considerable violence. 

The hydrochloric acid gas for this experiment must be 
dried, either by bubbling through strong sulphuric acid 
or by passing through a drying tube containing bits of 
porcelain or pumice stone moistened with sulphuric acid. 
The gas may be generated by the reaction of common 
salt with dilute sulphuric acid, four parts of water to 
about five of acid. If dried by passing through sulphuric 
acid, the rapidity of evolution of gas may be observed and 
regulated by increasing or decreasing the amount of heat 
applied. 

It is better, if possible, to collect the gas over mercury 
rather than by downward displacement, for in this way 
it may be obtained free from air. 



APPENDIX B 



353 



For this experiment, some straight graduated tube 
should be used, such as the eudiometer shown in some 
of the illustrations for the synthesis of gases. If this 
is not to be had, you ma) 7 use a burette, the capacity of 
which, both above and below the graduations, is accurately 
known. (See Fig. 67 for the general arrangement of the 




or 

^imiiiiiiiiiiiniiimiiimimi'iHin" 1 



Fig. 67 



apparatus.) When the graduated tube is completely filled 
with gas, put around it, as near the mouth as possible, a 
paper test-tube holder. This is made by folding a sheet 
of paper into a strip about an inch wide ; for use it is 
simply placed around the tube as shown in the accom- 
panying figure at iV, and grasped tightly between the 
thumb and fingers. The paper, being a 
poor conductor of heat, serves to prevent 
the transmission of the warmth of the 
hand to the glass so as to expand the 
hydrochloric acid. 

By means of this holder seize the tube, 
hold the thumb firmly over its mouth, 
and place in an upright position. Next, fig, 68. 




354 MODERN CHEMISTRY 

quickly drop into the tube a few grams of the sodium 
amalgam already prepared, and instantly replace the 
thumb, holding it as tightly as possible. Tip the tube 
back and forth a few times to hasten the action, and when 
this seems complete, place the mouth of the tube beneath 
the surface of the mercury and remove the thumb. The 
mercury instantly rises in the tube to fill the space 
formerly occupied by the chlorine, but now existing in 
the form of solid sodium chloride. Measure accurately 
the amount of gas remaining, and compare with the 
capacity of the tube ; what are your conclusions ? Test 
the residual gas with a light ; what is it ? What evidence 
have you that common salt is formed ? 

If the student finds he cannot hold his thumb tightly 
enough over the mouth of the tube to prevent leakage, 
he may use a short rubber stopper instead, and after the 
reaction of the sodium with the gas the tube may be 
opened crver water. 

6. Analytic Proof of the Composition of Hydrochloric 
Acid. — This may be furnished by the electrolysis of 
hydrochloric acid and the measurement of the gases 
obtained. Let the student arrange his own apparatus, 
and, taking such precautions as are necessary to avoid 
possible errors (mentioned in describing certain forms 
of electrolytic apparatus), make the experiment, and note 
results. 



APPENDIX C 

LABORATORY SUGGESTIONS 

1. Neatness. — To the best success in any chemical 
experiment neatness is absolutely essential; indeed, the 
merest traces of substances foreign to those with which we 
are working may cause a complete failure of the experi- 
ment. A student hardly knows what neatness is until he 
has had a thorough training in chemical analysis. 

The apparatus should always be clean when put away, 
and then before using should be rinsed with pure water. 
Never lay a cork or stopper down upon the table, as it will 
gather dust and thus pollute the reagent. If you desire 
to use some solution contained in *^rrm?*^ 

a bottle, take the stopper between 4^^fc^^ Ibf? 

the first and second fingers with the ^^^^^^^^B^ 
palm of the hand upward and ^^^IP?^ 

remove it from the bottle ; then ^%. 1 t\ 

without laying it down seize the (fft ||| 

bottle with the thumb on one side ;i 

and the fingers on the other. In 

this way the stopper will not come in contact with the 
side of the bottle and soil it, neither will dust and dirt 
be gathered from the table. The reagent bottles should 
be frequently wiped, as they soon become more or less 
covered with deposits which form from the gases gener- 
ated in the laboratory. The table also should be kept 
clean, and water and other liquids should not be allowed 
to remain if accidentally spilled. 

355 



356 



MODERN CHEMISTRY 



2. Order. — Great advantage will also be secured by 
having everything in its allotted place. Especially is this 
true of the reagent bottles, and the more there are of 
these the more important it is that they should be kept 
in order. For the larger schools probably about twenty 
reagent bottles will be furnished each student, and these 
will be arranged upon two shelves, one above the other. 
In such case, the following order is suggested as being as 
good as any : — 

Lower Shelf 

Beginning at left hand : — 
Sulphuric Acid . . Hydric Sulphate 



Hydrochloric Acid 
Nitric Acid . 
Ammonium Hydroxide 
Ammonium Chloride 
Ammonium Sulphide 
Ammonium Carbonate 
Barium Chloride . 
Potassium Dichromate 
Potassium Ferrocyanide 



Hydric Chloride 
Hydric Nitrate 
or Hydrate 



Potass. Acid Chromate 



H 2 S0 4 

HC1 

HN0 3 

NH 4 OH 

NH 4 C1 

(NH 4 ) 9 S 

(NH 4 ) 2 C0 3 

BaCl 2 

K 9 Cr 9 7 

K>eCy 6 



Upper Shelf 



Calcium Hydroxide or Hydrate . 

Mercuric Chloride 

Silver Nitrate 

Ferric Chloride 

Acetic Acid . 

Lead Acetate 

Potassium Iodide 

Sodium Carbonate 

Borax, powdered 



Argentic Nitrate 



Hydric Acetate 
Plumbic Acetate 



Crystals or powder 



Ca(OH) 2 

HgCl 2 

AgN0 3 

Fe 2 Cl 6 

HC 9 H 3 2 

Pb(C 2 H 3 9 

KI 

Na 2 C0 3 

Na 2 B 4 7 






) 2 



Some of the above reagents are known by different 
names, and in such cases two of them, the most common, 
have been given above. 






APPENDIX C 357 

3. Apparatus needed. — Each student should be assigned 
a locker where he may safely keep the apparatus supplied 
to him, and for the care of this he should be held respon- 
sible. The following apparatus is suggested : — 

3 Test-tubes, 5 x -|. 1 Test-tube Brush. 

3 Test-tubes, 6 x J. 1 Pair Forceps. 

3 Test-tubes, 6 x f . 1 Glass Stirring Rod. 

1 Evaporating Dish, small. 1 Blowpipe. 

1 Evaporating Dish, medium. 1 Platinum Wire. 

1 Beaker, 2 oz. 1 Rubber Cork, one hole. 

1 Small Flask, 2\ oz. 1 Rubber Cork, two holes. 

1 Delivery Tube. 1 Small Mortar. 

Directions will be given later for preparing the delivery 
tube, stirring rod, and some other desirable apparatus. 

The student should also have the following, and will 
furnish them himself : — 

An apron, reaching to the ankles. This may be made 
of denim, oil cloth, or rubber cloth. The last is the most 
serviceable in many ways, but is the most expensive. 
A Towel. An Iron Spoon. 

A Bar of Soap. A Clay Pipe. 

A Small Magnet. A Candle. 

A Small Triangular File. 

The candle will be needed very frequently during the 
first half of the work in studying the properties of gases. 
Common Property. — In addition to the individual prop- 
erty assigned above, certain articles on account of their 
size or for other reasons are used in common. There 
should be enough of them so that each member of the class 
may be supplied. Among these may be named : — 

An Iron Pan, 8 x 14 and about 1\ inches deep, to be 
used for a pneumatic trough. 



358 MODERN CHEMISTRY 

Test-tube Rack. Iron Ring-stand. 

Bunsen Burner, with Con- Funnel. 

nections. Wash-bottle (?). 

Wire Gauze. Sancl Bath. 

Manipulations 

4. Cutting Glass. — To cut tubing, with a sharp-cornered 
file scratch the glass entirely around where you desire to 
cut it. Now grasp the tube with both hands, the fingers 
above, and the thumbs below nearly meeting at the line 
scratched by the file. Now bend the tube downward and 
pull strongly apart at the same time. With a little prac- 
tice good square cuts may be made. The rough ends thus 
secured will cut any rubber connections used. To prevent 
this hold them in the Bunsen flame until the glass by 
becoming softened loses its sharp edges. 

Sometimes it becomes necessary to cut a bottle or large 
tube in two ; this may be done in two ways, but both de- 
pend upon the unequal heating of the glass. Tie around 
a bottle where you desire to cut it an ordinary twine 
string ; saturate it with kerosene and ignite it. Some- 
times it will be found necessary to apply the oil the second 
time, as soon as the first has ceased to burn, and again ignite 
it. In this way, if the oil has been applied carefully, a nar- 
row line extending around the bottle is heated strongly, 
and if the glass be cooled suddenly by pouring over it cold 
water, the bottle will be neatly severed. 

5. To prepare a Delivery Tube. — This may be made of 
rubber and glass tubing, or of glass alone. The former is 
often preferable because it allows of more freedom in 
manipulation. If made entirely of glass, two bends are 
necessary, and one should be within an inch of the end. 
Hold the tubing in the Bunsen burner, moving it back 



APPENDIX C 



359 




and forth and rolling it around so as to warm all por- 
tions equally. When the glass begins to soften, allow 
its own weight to bend it, and take 
care that you do not form a right- 
angled tube, but one of a gentle 
curve like the elbow of a stove pipe. 
When the bend has cooled just a 
little, close the openings at the bot- 
tom of the burner and hold the glass 
in the luminous flame until it is well 
covered with soot. This will cause 
the glass to cool slowly and hence 
make it less liable to fracture. Com- 
plete by making the second bend in 
the same way, forming an obtuse angle as shown in the 
figure. If rubber connections are used, a second bend 
is unnecessary. 

6. To make a Jet. — Frequently a tube drawn to a fine 
point is desirable. Take a piece of glass tubing 5 or 
6 inches in length and heat as in making a delivery tube. 
When it begins to soften, draw it .slowly apart until a 
tube of small diameter is obtained at the center, as shown 

in a in the adjoining 
figure. When some- 
what cooled, cut in two 
at a ; then make a bend 
in one of the shorter 
tubes, as shown in b. 
Round off the sharp edges and anneal as previously de- 
scribed. You will now have two jets, one straight and 
one bent, for both of which you will find uses. 

7. To make a Wash-bottle. — Any good-sized bottle or 
flask will do for this. The tube, a, should be drawn to a 




Fig. 71. 




360 MODERN CHEMISTRY 

jet as shown in the figure, and after being bent should 
reach nearly to the bottom of the flask. The other tube 
after being bent should just reach through the cork. By 
blowing through 5, a jet of water may be 
directed wherever desired; or if a larger 
stream is desired, it may be poured out at b. 
The bottle is more convenient if the jet, a, is 
attached to the rest of the tube by a rubber 
tube 2 or 3 inches long ; the stream of water 
may then be turned readily in any direction. 
The wash-bottle is indispensable for qual- 
itative work in washing precipitates. A 
rubber band should be slipped over the 
lower end of the tube, <x, so that if it strikes the side of 
the flask in removing the cork and tubing it will not 
be broken. If the lockers are too small to receive the 
wash-bottle, one may be used in common by the students 
working at each laboratory table or section. In such case 
it is better for each student to have a short tube with rub- 
ber connections to attach to 6, whenever he desires to use 
the bottle. 

8. To repair a Test-tube. — Test-tubes are frequently 
broken by the beginner, but they may be easily mended, 
and will then be almost as useful as at first. Hold the 
broken end in a hot Bunsen burner flame, roll the tube 
about to heat all sides evenly. When the glass becomes 
soft, by means of a glass rod, which will cohere to the 
softened tube, draw off the viscous portion, and thus seal 
the tube. Usually a small mass of softened glass will re- 
main upon the end of the tube. This must be drawn off 
in the same way, until the bottom is very thin, like the 
rest of the tube. Then by alternately heating and blow- 
ing into the tube, it may be rounded out and made almost 



APPENDIX C 



361 



as perfect as a new tube. After a little practice students 
may become skillful at this work. 

9. Blowpipe Work. — In metallurgy, the blowpipe must 
be used frequently, and two kinds of flames are employed, 
the reducing and th'e oxidizing. In preparing for either 
one, turn down the jet to about a quarter its usual force, 
or until you have a flame not much larger than that of a 
good-sized candle, and close the openings at the bottom so 
as to render it luminous. In the figure, a shows the small 
luminous flame ready for the use of the blowpipe, b shows 
the reducing flame. The tip of the blowpipe is placed 




Fig. 73. 



in the outer edge of the flame, and a gentle but steady 
stream of air forced into the flame. In this way a small 
luminous cone, I, will remain in about the center of the flame, 
and in this the metallic oxide should be held. This lumi- 
nous portion contains red-hot particles of carbon, and they 
have the power of reducing oxides of metals to the metallic 
condition. If this luminous cone is not apparent, too much 
air is being forced into the gas. Either blow more gently, 
or turn the gas on a little stronger. With a little practice 
the student will learn to breathe and blow at the same time, 
and will not find the work especially tiresome. 

For the oxidizing flame, c, above, the tip of the blowpipe 
is placed in the very center of the jet. In this way the 
air introduced and the gas become thoroughly mixed, and 
complete combustion ensues. The cone should be perfectly 



362 



MODERN CHEMISTRY 



non-luminous, and the metal to be oxidized should be held 
about where n is in the cut. The flame is exceedingly hot, 
and having an excess of oxygen readily reduces to oxides 
such metals as are oxidizable. 

10. Collecting Gases. — There are several methods for 
collecting gases, varying according to the characters of the 

gases. Those which are 
insoluble in water are usu- 
ally collected over water. 
Students will find an or- 
dinary baking pan, 2 
inches deep and about 6 
inches broad by 12 long, 
suiflciently large. The 
bottle to receive the gas 
is first filled with water 




Fig. 74. — Collecting over water. 



and inverted over the pan, Fig. 74. This is done by 
holding tightly a sheet of paper or glass over the mouth 
of the bottle until inverted and placed under the water 
in the pan. The delivery tube, T, dips under the bottle 
and conducts the gas from the generating flask, Gr, into the 
bottle. 

11. Collecting by Downward 
Displacement. — Gases soluble in 
water obviously cannot be col- 
lected by the method already 
described. If it is necessary 
to have them absolutely pure, 
mercury is frequently substi- 
tuted for the water. Ordina- 
rily, however, if heavier than air 

they are collected by doivnward displacement. By this 
method the bottle is simply left standing upon the table, 




Fig. 75. 



APPENDIX G 



363 



and the delivery tube reaches down into the bottle. 
Thus the heavier gas is introduced below the air, and 
gradually displaces it. Such gases as chlorine or carbon 
dioxide are collected in this way. If the gas is lighter 
than air and soluble in water, it is usually collected by 
upward displacement. The receiving bottle is held in an 
inverted position, and the delivery tube runs up to the 
bottom of the bottle, gradually displacing the air in the 
bottle. In Fig. 75, a shows the arrangement for collect- 
ing by dowmvard displacement, and 5, that for uptvard 
displacement. 

12. Measurements. — Frequent reference is made through- 
out this work to the cubic centimeter and gram, and tire 
student should have fairly definite ideas 
of these terms. This can come only by 
practice. For the volumetric, a test-tube 
and beaker may be graduated. From a 
burette run into a test-tube 1 cc. of water ; 
indicate its height by fastening upon the 
tube just above the lowest part of the 
meniscus a narrow strip of mucilage paper. 
Add another cubic centimeter and mark 
the height in the same way. Thus grad- 
uate - the tube up to 5 cc. ; mark it also 
for the 10 cc. Now that the graduation l cc ; 
may be permanent, with a file scratch 
carefully the marks, after which the paper 
may be removed ; a shows the meniscus 
for each cubic centimeter, and b the small strip of paper. 
In the same way graduate a beaker for 5, 10, 15, 20, and 
25 cc. 

As different compounds vary so greatly in density, it is 
more difficult to obtain an accurate idea of a gram, but the 




Fig. 76. 



364 MODERN CHEMISTRY 

student should be able to approximate it. Put upon one 
scale pan of a balance a small evaporating dish, and coun- 
terbalance it with shot or sand upon the other. Then add 
a gram weight to the shot. Into the evaporating dish now 
slowly add common salt until the gram weight is balanced. 
Thus try some other amount, as 2 g. or 5 g. 

If the classes are large, one portion may be graduating 
the test-tubes and beakers, while another is doing the 
gravimetric work. This will greatly expedite matters. 

13. Precipitates. — A precipitate is any solid matter 
thrown down in a solution by adding to it some reagent. 
It may be very dense, so as to be quite jelly-like, or it may 
form merety a cloud in the solution. To illustrate, put 
one drop of sulphuric acid into a beaker half or two-thirds 
full of water and add 2 or 3 cc. of barium chloride solution. 
The dilute solution should thus give a slight precipitate 
only. Now powder about a gram of ferrous sulphate and 
dissolve in as little water as possible, 2 or 3 cc, then add 
a few drops of ammonia. A thick gelatinous precipitate 
should form. 

14. Decanting and Filtering. — These are processes for 
separating a precipitate from the solution in which it is 
formed. When the precipitate is one that has considerable 
density and settles quickly, leaving a clear solution, this 
supernatant liquid may be decanted or poured off. There 
is no objection to this method unless the presence of small 
particles of the precipitate in the decanted portion, or of 
the solution in the precipitate, will interfere with subse- 
quent tests. To illustrate, a solution of lead acetate may 
be precipitated with hydrochloric acid, and after warming 
slightly and allowing the precipitate to settle, the solution 
may be decanted. 

But in cases where the separation must be complete, 



APPENDIX C 



365 



filtration is necessary, that is, passing the solution through 
a filter paper. There are two ways of folding filters : the 
simplest, and one used 
when the precipitate is to 
be removed from the pa- 
per, is as follows : fold the 
paper to form a semicircle, 
5, then this to form a 
quadrant, making one fold 
slightly smaller than the other 




Folded Twice Opened 




This is done because 
funnels are seldom perfectly made, 
and one " quarter" will fit them 
better than another. Usually this 
is the larger. Now open out one 
of the quarters, and press down 
neatly into the funnel. If the 

quarter tried does not seem to fit, the other one may do 

so better. Now moisten with a little water, and with the 

fingers press the paper 

against the sides of the 

funnel to remove any air 

bubbles that may exist 

there. In filtering, pour 

in slowly at first, especially 

if the precipitate is very 

finely divided. If the solu- 
tion does not come through 

clear, it may be necessary 

to filter again through the 

same filter paper. The 

pores will soon become 

partially filled, and the fil- 
trate will be perfectly clear. Fig. 79. 




366 MODERN CHEMISTRY 

In filtering, the stem of the funnel should always be made 
to touch the side of the beaker or vessel into which the 
liquid is being passed, so that no drops may spatter out. 
Furthermore, in pouring a liquid from any vessel, it 
should always be allowed to run down a moistened stir- 
ring rod into the funnel. By observing these precautions, 
neatness in transferring liquids from one vessel to another 
will be secured. 

15. Opening Bottles. — The common acids and aqua 
ammonia, as well as some other reagents, are frequently 
put up in bottles with glass stoppers. They are sealed by 
dipping the stopper into melted paraffin before inserting 
into the bottle. To remove the stopper the paraffin must 
be melted. This may be done by turning down the gas- 
jet moderately low, taking the bottle in both hands, hold- 
ing the neck over the flame, not too close, and rolling it 
rapidly around so as to heat all sides alike. Be careful to 
heat the glass only gently. In a moment or two the wax 
will be melted and the stopper may be very easily removed. 
With a little practice bottles may be opened in this way 
without ever breaking or cracking. Be careful, however, 
in removing the stopper, never to have the face directly 
over the bottle. 

16. Platinum Wires. — These are used in making flame 
and borax-bead tests for various metals. For the sake of 
convenience in handling, they are generally fused into a 

short piece of glass tub- 
Take a few inches 
of small-size tubing and 
FlG ' 80 * draw out, as in making 

a jet such as has already been described for use in testing 
the combustibility of gases. Cut the glass in two, as 
before, and insert the platinum wire into the tubing to 




APPENDIX C 



367 



mrnnr 



a distance of 3 or 4 cm. ; again hold in the flame until 
the glass is softened. Upon cooling, the wire will be 
securely fastened in the tubing. (See Fig. 80.) 

17. Electrolytic Apparatus. — If necessary, the student 
may prepare his own apparatus for experiments in electrol- 
ysis out of other apparatus that 
he will find at hand. Take two 
pieces of heavy platinum wire, 
each about a foot long, and make 
into spiral coils by wrapping 
around a pencil. Leave two or 
three inches straight at one end, 
as shown at «, Fig. 81. 

Fit to a short-necked bell jar 
with an open top a rubber cork 
with two small holes, and support 
the bell jar upon an iron ring- 
stand, fastening it securely in 
position. Next, take two pieces 
of small glass tubing, each long 
enough to reach through the cork 
<?, and extend just into the body of the jar. Insert the 
straight ends of the two platinum spirals, already made, 
through these tubes, and fuse the glass at the ends so as to 
fasten the wires firmly in the glass ; make a small loop in 
the wire at the lower end. See b in the figure. Insert 
the two electrodes thus prepared through the holes in the 
cork, and see that everything is water tight. 

Next take two burettes with glass stop-cocks and deter- 
mine accurately the capacity of each below the point of 
graduation, that is, from m to n in Fig. 82. This must 
be done if we desire to measure accurately the amount 
of gas collected. Now by means of clamps support these 




Fig. 81. 



368 



MODERN CHEMISTRY 




Fig. 82. 



two burettes inverted over the two spiral electrodes, and 
the apparatus is complete. For use, fill the bell jar with 
the liquid to be electrolyzed to some distance 
above the mouth of the burettes. Attach a 
rubber tube to the tip of the burettes, open the 
stop-cock, and by suction fill each with the liquid 
and close the stop-cock. Turn on the current, 
and the capacity of each burette above the point 
of graduation having been determined, the 
amount of gas which collects in each tube is 
quickly read. 

Instead of the burettes, test-tubes 8 inches by one-half 
in diameter may be used with good results, except that 
the gases cannot be accurately measured. 

18. A Simple Electrolytic Apparatus. — Occasionally it 
may be desired to electrolyze a substance without sepa- 
rating the gaseous products. For such purposes a very 
simple form of apparatus may be employed, as shown in 
the figure. Prepare the two electrodes as 
described for the more complicated form, 
and fit them to a 3-hole stopper as 
shown in Fig. 83. Through the other 
opening pass a bent delivery tube, T, for 
conducting off the mixed gases which will 
collect in the top of the bottle when the 
current is passed. 

Such apparatus as this may be used to 
show the explosive character of the mix- 
ture of hydrogen and oxygen obtained by 
the electrolysis of water, or of hydrogen and chlorine 
resulting from the decomposition of hydrochloric acid. 
To prevent the contents of the bottle becoming too warm, 
it should be placed in a vessel of cold water. Use hydro- 




Fig. 83. 



APPENDIX C 



369 



chloric acid of specific gravity about 1.1, and allow the 
current to pass for some time before collecting the gases, 
in order that the liquid may become saturated with the 
chlorine. If it is desired to collect bottles of the mixed 
gases over water, let the water be first saturated with 
common salt. 

19. Eudiometers. — The eudiometer is an instrument 
used to test the composition of mixed gases. The most con- 
venient form for all purposes is the U-shaped one, in which 
mercury is used to confine the gases. The air left in one 
limb of the tube 
serves as an air cush- 
ion to receive the 
shock of the explo- 
sion. The straight 
eudiometer, how- 
ever, is cheaper, and 
with a few addi- 
tional attachments 
may be used satis- 
factorily. A in the 
figure is an open-top 
bell jar, such as has 
been used in other 
experiments. The 
neck of A is closed with a tight-fitting, 1-hole rubber 
stopper, through which passes a glass tube having an en- 
largement blown upon the lower end, at B. Another 
rubber cork, which must fit the eudiometer, E, very tightly, 
is put upon the glass tube as shown in the figure. This 
must also fit very tightly. T is simply a piece of glass 
tubing about one inch in diameter, which should have a 
capacity somewhat greater than E. It is closed at the 




Fig. 84. 



370 MODERN CHEMISTRY 

lower end with a cork, through which passes a short glass 
tube. A rubber tube connects the two portions of the 
apparatus, and just above B is fastened by some fine insu- 
lated copper wire wrapped about it. 

For use the eudiometer is filled with water and sup- 
ported in position over A. The gases to be exploded are 
introduced separately, and each measured carefully, the 
eudiometer being held by a paper test-tube holder at such 
height that the water stands at the same level inside and 
outside. Now press JE firmly down upon its cork, and 
lower T as much as possible in order that the confined 
gases may have the pressure upon them reduced ; grasp 
the rubber tubing near B firmly with the thumb and 
finger, and pass the spark. After the explosion, adjust 
the level inside and outside of JE as when the gases were 
introduced, and measure the residue. If this adjustment 
cannot be secured by lowering _Z£, it may remain connected 
as when the spark was passed, and the level secured by 
changing the height of T. 

20. Aspirators and Aspirating Bottles. — As an aid in 
filtering certain classes of precipitates, an aspirator is fre- 
quently used. This acts upon the principle 

4 of the Sprengel air-pump. The aspirator con- 

sj^ „ > sists merely of two tubes, A and B, secured 

"** jj at right angles to each other. A is attached 
t| to a water faucet, and B, by means of heavy- 

y walled rubber tubing, to a filter flask. As 

Fig. 85. .Q ie wa ter flows through A, the air is gradu- 
ally withdrawn from the flask; the pressure being thus 
removed from beneath the filter containing the precipitate, 
the liquid is forced through much more rapidly. 

The filter flask is usually shaped like an Erlenmeyer 
flask (see Fig. 86), and has a side tube for connecting 



APPENDIX C 



371 




Fig. 87. 




Fig. 86/ 



with the aspirator at B. It is made of heavy glass so 
as to withstand any ordinary atmospheric pressure. 
For use it is fitted with a rubber 
stopper haying one hole, through 
which the stem of a funnel is inserted. 
In the apex of the funnel is 
placed a small platinum cone, 
perforated with minute open- 
ings. This cone is used to 
prevent the breaking of the 
filter paper by the atmospheric pres- 
sure; at the same time the numerous 
small holes permit the outflow of the 
filtrate with comparative freedom. 

For certain experiments an aspirat- 
ing bottle is almost indispensable. For example, suppose 
the experimenter desires to cause a regular flow of air or 
of some other gas through some vessel, suitable apparatus 
is necessary and may be very easily made. Large bottles, 

holding 3 or 4 liters, 
will serve best. To 
each fit a cork with 
two holes, and insert 
glass tubing as shown 
in the accompanying 
figure. The bent 
tube, (7, has attached 
a short piece of flex- 
ible rubber tubing, 
upon which is placed 
a screw clamp, at H. 
By means of this the 
Fig. 88. flow of gas issuing 




372 MODERN CHEMISTRY 

from N is regulated. The bottle, M, is placed upon a 
box so as to elevate it considerably above N. A rubber 
tube, U, connects the two bottles, and, being flexible, 
allows of the elevation of either bottle above the other. 

If you desire to fill JV with any gas not soluble in water, 
place both down upon the table, and fill JV* completely 
with water. Open the clamp at H, and insert the cork 
with the tubing into N. The water will be forced out 
into Gr, and expel the air therefrom ; this done, connect 
at J?" with the generating flask (not shown in the figure), 
after having waited until all air has been expelled from it. 
By the gas pressure, the water will be forced from N over 
into M ; continue until N is nearly filled, close the clamp 
at R tightly, and remove the generator. Elevate M to its 
position upon the box, and the aspirator is ready for use. 

By simply opening the screw clamp, the siphon connect- 
ing the two bottles transfers the water from M to N as 
rapidly as the exit of gas at H will allow. If the gas has 
been permitted to fill completely the bottle iV, and has 
forced the water out of the siphon tube, it is only neces- 
sary to apply a little pressure at D. If a dry gas is de- 
sired, it must be obtained by passage from IV through 
some suitable drying tube attached at IT. If the gas to 
be used is ordinary air, the action of this apparatus may 
be made continuous, except for a momentary delay in 
changing the connections, by placing first M, and then iV, 
upon the box, and connecting the receiver with the tubes, 
Gr and D, respectively. 

The apparatus may be used in this way for showing the 
presence of carbon dioxide in air, by forcing it through 
lime-water. In other cases, where the amount of gas 
needed is not in excess of the capacity of the bottle N, 
this apparatus will work with entire satisfaction. 



APPENDIX G 



373 



21. Gas Generators. — It is often desirable to have a 
generator, automatic in action, which will furnish a steady 
flow of gas and be ready for use at a moment's notice. 
Kipp's apparatus meets such a demand ; but at much less 
expense one which works equally well may be prepared 
for any laboratory. In the fig- 
ure, A is a bottle of about 500 cc. 
capacity, fitted with a cork and 
tube at P, to keep out dust. 
Through the bottom at if, with 
a glass drill, make a hole and 
insert a rubber cork with one 
perforation. 

Through B near the bottom 
drill a hole and insert a rubber 
cork with a glass tube and short 
rubber connection clamped with 
a Hoffman screw. This is for 
the purpose of drawing off the 
spent acid. In the top of B fit a 
stopper with two holes ; through 
one of these pass a long tube 
reaching to the bottom of B and 
extending up into A. To the other hole fit the bent tube, 
D, which has rubber connections for joining with any 
other apparatus. When not in use, this is kept tightly 
closed with a screw clamp. 

If you desire to use this apparatus as a hydrogen gener- 
ator, place a half pound or more of zinc in B, close tightly 
the screw clamp at i), and pour diluted sulphuric or hydro- 
chloric acid into A until about two- thirds full. Open the 
screw clamp ; the acid will run down into the lower bottle 
and will continue to react with the zinc as long as the gas 




Fig. 89. 



3f4 MODERN CHEMISTRY 

has free exit at D. If, however, the clamp is closed, the 
pressure in B soon becomes sufficient to force the acid up 
the longer tube into the upper bottle, and the evolution 
of gas ceases. 

The bottle, A, is held in position by a clamp at the neck, 
and rests upon a ring of the support. The holes at iT and 
JE may be drilled by using a large file broken off, together 
with emery dust. To use the generator for hydrogen 
sulphide or carbon dioxide, the zinc would be replaced 
with ferrous sulphide or marble. 

22. Correction of Barometric Reading. — In the various 
problems given in the text in connection with the Law of 
Charles, it was assumed without being stated that we were 
dealing with dry gases. Further than this, in the quanti- 
tative work with gases, certain corrections have been neg- 
lected. For exact work, however, in the measurement of 
gases, not only must the temperature be known, and the 
barometric pressure as well, but also certain other facts. 
If the gas has been collected over water, the exact volume 
will not be obtained by methods already used, for the 
reason that the presence of water vapor increases the 
tension of -the gas, and hence the volume. In reducing 
the volume of gases, therefore, to standard conditions, 
allowance must be made for this tension. This has been 
carefully estimated, and for the ordinary range of tempera- 
ture is shown below : — 



19° C. . 


. 16.35 mm. 


22.0° C. . 


. 19.66 mm 


19.5° C. . 


. 16.86 " 


22.5° C. . 


. 20.27 " 


20.0° C. . 


. 17.39 « 


23.0° C. . 


. 20.89 " 


20.5° C. . 


. 17.94 " 


23.5° C. . 


. 21.53 " 


21.0° C. . 


. 18.50 « 


24.0° C. . 


. 22.18 " 


21.5° C. . 


. 19.07 « 


24.5° C. . 


. 22.86 " 



APPENDIX C 375 

To illustrate, suppose we have 40 cc. of gas, the tern- 
perature of the room being 21° C, the barometric pressure 
740. According to the law, stated previously, — 

TZ V 1 X P' 
or T= , 

in which V represents volume under standard pressure P, 
which is 760, V the given volume of gas under the pres- 
sure P f . Substituting, — 

F== 40 x 740 
760 

But making correction for aqueous tension, we have 

\ - p ' 

in which p is the tension of the aqueous vapor. From the 
table given above, we find that at 21° C. this is 18.5 mm. 
Substituting in the formula, we have, — 

r= 40 x (740-18.5) 

760 

which will give the true volume of the gas under standard 
conditions. 

23. Drying Tubes. — Drying may usually be accom- 
plished by forcing a strong current of air through the 
tube by means of a foot-bellows , if the tube has been 
previously moistened with alcohol, the process will be 
materially hastened. In like manner flasks may be dried. 
By means of rubber tubing connect a glass tube, long 
enough to reach to the bottom of the flask, to a foot-bel- 
lows, and direct a strong current of air into the flask, 



376 MODERN CHEMISTRY 

24. Recording Results of Experiments. — In the first 
place, the student should understand exactly what he is 
expected to learn from the experiment ; then he must 
know what steps are necessary in order to secure the 
correct results. Do not make the mistake of drawing 
conclusions before the experiment is complete, and then 
endeavoring to make the results conform to your precon- 
ceived ideas. Learn to see everything that occurs, and 
draw your conclusions in accordance with what really 
happens. 

These results should be recorded in suitable note-books, 
and, were it possible, always completed in the laboratory. 
Note the results neatly and concisely in good rhetorical 
sentences. When they admit of being tabulated, such a 
form is always desirable. If the notes are not written up 
in the laboratory, a brief record should be made there, 
and at home put into permanent form in the note-book 
without delay. These records should be examined fre- 
quently by the teacher, at least after the completion of 
each distinctive portion of the work; for instance, in 
studying the halogen group, when the work in chlorine 
has been done, the notes should be examined ; after that 
in bromine is completed, a similar examination should 

take place. 

Preparing Solutions 

25. For ordinary work, reagents which are "commer- 
cially pure " will do, and are much cheaper. It is better 
to use distilled water in making up all solutions, but for 
some, such as caustic potash, soda, and such as form pre- 
cipitates with water that is more or less "hard," pure 
water is essential. 

26. Acids — Hydrochloric, Nitric, and Sulphuric. — For 
ordinary work these acids should be diluted with twice 



APPENDIX C 377 

their own volume of water. In the case of the last acid 
the water must be added very cautiously, as great heat is 
generated. It is better to take what water is to be used 
in diluting the acid, and very gradually add the sulphuric 
acid to it. Acetic acid may also be diluted. When an 
acid stronger than the one prepared in this way is de- 
manded, it is so stated in the text. 

27. Ammonia. — Ordinary aqua ammonia should be 
diluted with about three parts of water. 

28. Ammonium Chloride. — This should be made up 
with about 100 g. of the salt to a liter of water. 

29. Ammonium Carbonate. — About 200 g. to liter. 

30. Ammonium Oxalate. — About 40 g. to liter. 

31. Ammonium Sulphide. — This may be prepared by 
the teacher if preferred. It is done by taking ammonium 
hydroxide as diluted above and passing into it a current 
of hydrogen sulphide until saturated. If yellow ammo- 
nium sulphide, (NH 4 ) 2 S X , is desired, add to the ammonia 
at the beginning a little sulphur in the form of flowers. 
When the solution is saturated, it is customary to add 
to it about two-thirds as much more of the ammonium 
hydroxide. 

32. Barium Chloride. — About 100 g. to the liter. 

33. Potassium Dichromate. — About 50 g. to the liter. 

34. Potassium Ferrocyanide. — About 75 g. to the 
liter. 

35. Calcium Hydroxide. — Saturated solution. 

36. Mercuric Chloride. — Saturated solution. 

37. Mercurous Nitrate. — About 50 g. to the liter, with 
about one-twentieth part of nitric acid added. Otherwise 
a basic salt forms in the solution. It is a very good plan 
to put a few drops of mercury into the bottle containing 
the solution. 



378 MODERN CHEMISTRY 

38. Silver Nitrate. — About 50 g. to the liter. Keep 
the solution in an amber-colored bottle and away from 
contact with organic substances. 

39. Ferric Chloride. — About 50 g. to the liter. 

40. Ferrous Sulphate. — This must be made up as 
desired. About 100 g. to the liter. 

41. Lead Acetate. — About 100 g. to the liter. 

42. Potassium Iodide. — About 50 g. to the liter. 

Other Solutions used Occasionally 

43. Arsenic Chloride. — Dissolve arsenious oxide, As 2 3 , 
in caustic soda, and then add hydrochloric acid until the 
solution gives an acid reaction. 

44. Antimony Chloride. — Add hydrochloric acid to 
water until well acidulated, and then a small quantity of 
antimony trichloride ; a solution of antimony may be 
obtained from the antimony tartrate in the same way. 

45. Bismuth Nitrate. — This must be prepared in the 
same manner as the antimony chloride. Dissolve a few 
crystals of the salt in water to which considerable nitric 
acid has been added. 

46. Calcium Chloride. — About 100 g. to the liter. 

47. Calcium Sulphate. — Saturated solution. 

48. Cobalt Nitrate. — About 50 g. to liter. 

49. Chromium Chloride. — Prepare as indicated in the 
text. To a solution of potassium dichromate add about 
one-twentieth as much hydrochloric acid and a little 
alcohol, and boil. The green solution obtained will be 
chromium chloride. 

50. Copper Sulphate. — About 50 g. to liter. 

51. Di-sodium Phosphate. — About 100 g. to liter. 

52. Potassium Cyanide. — About 100 g. to liter. 

53. Potassium Chromate. — About 50 g. to liter. 



APPENDIX C 



379 



54. Potassium Hydroxide. — About 100 g. to liter. 

55. Sodium Hydroxide. — About 100 g. to liter. 

56. Magnesium Sulphate. — About 100 g. to liter. 

57. Sodium Carbonate. — About 100 g. to liter. 

58. Lead Nitrate. — About 100 g. to liter. 

59. Stannous Chloride. — First acid about one-twentieth 
part of hydrochloric acid to the water, and then about 
75 g. of the solid to a liter. It is better to put a piece of 
granulated tin into the solution. 

60. Cochineal Solution. — Grind up the solid in a mortar 
and dissolve in water or in a 10 per cent solution of alcohol. 

61. Indigo Solution. — Treat about 1 g. of indigo with 
about 10 g. of sulphuric acid. After standing several 
days, dissolve the whole in water. 

62. Litmus Solution. — Dissolve the blue solid, powdered, 
in water. 

63. Phenol -phthalein. — Dissolve about 1 g. in 100 cc. 
of 50 per cent alcohol. 

64. Ammonium Molybdate. — Dissolve 15 g. of am- 
monium molybdate crystals in 100 cc. of aqua ammonia 
as prepared above. To this add an equal volume of 
distilled water, and finally 125 cc. of nitric acid, specific 
gravity about 1.4. 

Supplies Needed 



65. Chemicals. — For ten students. 

Acid, Acetic 1 lb. 

" Hydrochloric . . . 10 " 

" Nitric 6 " 

" Oxalic J " 

" Sulphuric . . . . 10 " 

" Tartaric J " 

Alcohol 1 qt. 

Alum lib. 



Ammonium, Carbonate . 


lib 


u 


Chloride 


1 " 


a 


Hydroxide . 


8 « 


a 


Nitrate . . 


l a 


a 


Sulphate . 


1 a 

8 


Antimony, 


Metallic . . 


1 u 
4 


a 


Potassium Tar- 






trate . . . 


i « 



380 



MODERN CHEMISTRY 



Arsenic, Metallic .... 

" Trioxide . . . . 

Barium Chloride .... 

Bismuth, Metallic .... 

" Nitrate . . . . 

Bleaching Powder .... 

Calcium, Carbide .... 

" Chloride .... 

" Fluoride .... 

" Sulphate. . . . 

Carbon Disulphide . . . 

Charcoal, Powdered, animal, 

Stick 

" Wood, powdered, 

Cobalt Nitrate 

Cochineal 

Copper, Metallic, turnings . 

" Nitrate 

" Sulphate . . . . 

Ether . . 

Indigo 

Iodine 

Iron, Filings 

" Chloride 

" Sulphate 

" Sulphide 

Lead, Metallic 

" Acetate 

" Nitrate 

" Oxide, Litharge . . 
" " Minium . . 

Litmus 

" Papers, Red and Blue, 

each 

Magnesium Ribbon . . . 

" Powdered . . 

" Sulphate . . 

Manganese Chloride . . . 

" Dioxide . . . 



Jib. 


4 


i " 


i " 


1 " 
4 


1 « 
4 


i " 


J " 


i " 


i " 


i " 


l u 


1 doz. 


lib. 


i " 


1 u 

8 


2 " 


1 u 

8 


J " 


1 U 
4 


16 


1 u 
16 


i " 


i " 


1 U 
2 


2 " 


i " 


i a 

4 


i " 


4 


\ " 


Alb. 


Jq- 


^ib. 


1 a 

4 


I " 


i " 


1 " 



Mercury, Metallic .... J lb. 

" Bichloride . . 

Mercuric Nitrate . . . 

" Oxide . . . 

Mercurous Nitrate . . 

Phosphorus, Ordinary . 

Potassium, Metallic . . 

" Bromide . . 

" Carbonate . 

" Chlorate . . 

" Chromate 

" Cyanide . . 

" Dichromate . 

" Ferrocyanide 

" Iodide . . . 

" Hydroxide, sticks, 1 " 

" Nitrate . . 

" Nitrite . . 

" Permanganate 

" Sulphocyanide 

Silver Nitrate .... 

Sodium, Metallic . . . 

" Borate (Borax) 

" Carbonate . . 

" Chloride . . . 

" Hydroxide, sticks 

" Phosphate, Di . 

" Sulphite . . . 

" Thiosulphate . 

Starch 

Strontium Nitrate . . 

Sugar 

Sulphur, roll 

" flowers . . . 

Tin, Metallic .... 

" Chloride .... 

Turpentine 

Zinc, Granulated . . . 
« Dust 



i 

T6 



i 
i 

i 






APPENDIX D 

REFERENCE LIBRARY 

No text on chemistry can hope to give more than a 
glimpse at the subject. Naturally, therefore, it should 
be the aim of every teacher to build up a reference library 
for the use of himself and students. Among the many 
good books to be obtained, the following are suggested : — 

Neivth's Inorganic Chemistry — Longmans. 

Neivth's Chemical Lecture Experiments — Longmans. 

Mendeleeff's Principles of Chemistry — Longmans. 

Ostwald^s Outlines of General Chemistry — Macmillan. 

Ostwald's Foundations of Analytical Chemistry — Mac- 
millan. 

Walker-Dobliri s Chemical Theory for Beginners — Mac- 
millan. 

Roscoe and Schorlemmer } s Treatise on Chemistry, Vols. I 
and II — Appleton. 

Remsen's Chemistry, Advanced Course — Holt. 

Remsen's Theoretical Chemistry — Lee. 

Ramsay s Experimental Proofs of Chemical Theory — 
Macmillan. 

Cornish's Practical Proofs of Chemical Laws — Longmans. 

Johnston's Chemistry of Common Life — Appleton. 

Lassar-Cohn's Chemistry of Every-day Life — Lippincott. 

Ramsay's Gases of the Atmosphere — Macmillan. 

Meyer's History of Chemistry — Macmillan. 

TJiorpe's Essays in Historical Chemistry — Macmillan. 

381 



382 MODERN CHEMISTRY 

Sutton's Volumetric Analysis — Blakiston. 

Addy man's Agricultural Analysis — Longmans. 

Alembic Club Reprints — Chemical Pub. Co., Easton, Pa. 

Foundations of the Atomic Theory. 

Experiments on Air. 

Foundations of the Molecular Theory. 

Discovery of Oxygen. 

Elementary Nature of Chlorine. 

Liquefaction of Gases. 

Early History of Chlorine. 
Muir's Heroes of Science — Young & Co. 
Shenstone's Glass Blowing — Longmans. 
Thorpe's Chemical Preparations — Ginn. 

APPENDIX E 

BIOGRAPHICAL 

The following are among those who have contributed to 
chemical literature or to the advancement of the science. 

Age of Alchemy 

Greber. — Arabian alchemist of eighth century ; author 
of several chemical works, and discoverer of aqua regia. 

Albertus Magnus. — Died 1280. Advanced the theory 
that the metals were composed of water, arsenic, and 
sulphur. 

Bacon, Roger. — - Thirteenth century. English alche- 
mist. Advocated experimental proof of chemical theory. 
Inventor of gunpowder. 

Valentine, Basil. — Fifteenth century. Wrote several 
works on chemistry. Probably a fictitious name of 
Johann Tholde. 






APPENDIX E 383 

Medical Era of Chemistry 

Paracelsus, a name coined for himself by Theophrastus 
Bombastus von Hohenheim. — Early part of the sixteenth 
century. By his study and preparation of a large number 
of medicines, he earned for himself the title, " Father of 
Medicine." 

Libavius. — Died in 1616. Proceeded with the work 
begun by Paracelsus. Wrote a Handbook of Chemistry. 

Van Helmont, Jean Baptiste. — 1577—1644. Discov- 
ered several gases. 

Boyle, Robert. — 1627-1691. Real founder of the 
sciences of physics and chemistry. Formulated Boyle's 
Law, and advanced the true theory as to the composi- 
tion of matter. 

Becher, Johann Joachim. — 1635-1682. German chem- 
ist. Author of theory that when a metal burns terra 
pinguis escapes from it. 

Age of Phlogiston 

Stahl, G-eorg Ernst. — 1660-1734. Founder of the phlo- 
gistic theory of combustion, that all combustible sub- 
stances contained an unknown something called phlogiston 
which escaped when the substance burned. It was an 
outgrowth of Becher's theory. 

Hoffmann, Christoph Ludivig. — 1721-1807. Physicist 
and chemist. His theory of the reduction of a metal was 
about the same as that held to-day. He believed that 
the calces of the metals contained the metals themselves 
and some other substance, wdiich he called sal acidum. 

Black, Joseph. — 1728-1799. Professor of chemistry in 
Edinburgh. Discovered carbon dioxide and proved that 



384 MODERN CBEMISTRY 

the carbonates of the alkalies and alkaline earths are 
not elements. 

Cavendish, Henry. — 1731-1810. Discovered hydrogen ; 
studied the composition of water and the air, and made 
a large number of experiments with the latter. Prepared 
nitric acid by synthesis. 

Priestley, Joseph. — 1733-1804. Discoverer of oxygen, 
and strong advocate of phlogistic theory. 

Scheele, Carl Wilhelm. — 1742-1786. Discoverer of 
chlorine; made some investigations in organic chemistry; 
prepared glycerine and prussic acid. 

Modern Era of Chemistry 

This coincides roughly with the nineteenth century. 

Lavoisier, Antoine Laurent. — 1743-1794. Founder of 
modern chemistry. Made a beginning in quantitative 
work, and overthrew the theory of phlogiston. Advanced 
the idea of the conservation of matter. 

Gcay-Lussac, Joseph Louis. — 1778-1850. Author of 
the law of combination of gases by volume. Made an 
extensive study of the general properties of gases ; deter- 
mined the relation between the volume of a gas and its 
temperature, thus supplementing Boyle's work. 

Berzelius, Johann Jacob, Baron. — 1779-1848. Studied 
the atomic weights of the elements ; improved the usual 
methods of chemical analysis, and investigated the law 
of combining proportions. 

Proust, Louis Joseph. — 1760-1826. Advocated the theory 
that the elements combine always in definite proportions, 
now known as the "Law of Definite Proportions." 

Dalton, John. — 1766-1844. Advanced the atomic 
theory of matter, and formulated the " Law of Multiple 
Proportions." 



APPENDIX E 385 

Berthollet, Claude Louis. — 1748-1822. Made a long 
series of experiments, studying the behavior of ammonia, 
hydrogen sulphide, chlorine, and other gases. 

Davy, Sir Humphry. — 1778-1829. Studied the prop- 
erties of various gases ; proved that the alkalies, caustic 
soda and potash, are not elements. 

Dulong and Petit. — Early part of nineteenth century. 
Made a study of the metals. Formulated the law that 
the specific heats of the metals are inversely proportional 
to their atomic weights. 

Dumas, Jean Baptiste Andre. — 1800-1881. Made an 
extensive study of vapor densities. 

Faraday, Michael. — 1791-1867. Succeeded in liquefy- 
ing many of the gases ; studied physical chemistry, and 
determined the effects of an electric current upon electro- 
lytes. He formulated the " Law of Definite Electrolytic 
Action,'' that an electric current decomposes electrolytes 
so that equivalent amounts of the substance are liberated 
at the kathode and anode. 

Liebig, Justus, Freiherr von. — 1803-1873. Studied 
organic chemistry ; investigated the phenomenon of 
isomerism. 

Mendeleeff, Dmitri Ivanovich. — Born 1834. Russian 
chemist. Formulated the " Periodic Law of the Ele- 
ments." Author of general chemistry. 

Pictet and Cailletet. — Physico-chemists of the present 
time. They have done much work in producing low 
temperatures, and in liquefying air, hydrogen, and 
oxygen. 

Ramsay, William. — Born 1852. Discoverer of argon 
in 1894. English scientist of to-day. 

Deivar, James. — Born 1842. English scientist of the 
present time. Has studied carefully low temperatures. 



386 



MODERN CHEMISTRY 



Moissan, Henry. — French chemist of the present time. 
Has succeeded in preparing artificial diamonds; has 
also studied carefully the properties of liquid fluorine. 



Meaning of Alchemistic Terms 

The student in attempting to read the reports of the chemists of 
the eighteenth century will find much difficulty in understanding the 
alchemistic terms so universally employed. The following are among 
those most commonly met with, and are given to encourage the student 
to read these accounts himself. The Alembic Club Reprints, men- 
tioned among the books suitable for reference, furnish the most 
desirable portions of the writings of such investigators as Scheele, 
Dalton, Priestley, and others. It will be noticed that often several 
terms are used for the same substance. This was in accordance w r ith 
the plans of alchemy to keep secret the discoveries and mystify any 
who might attempt to decipher the records. 



Old Terms 

Acid . 

Acid of chalk 

Acidum salis 

Aer fixus . 

Air 

Alkali of tartar 

Aqua fortis 

Aqua regis . 

Azotic gas . 

Blanc d'Espagne 

Calx . 

Calx of silver 

Colcothar . 

Dephlogisticated air 

Draco mitigatus 

Fire air 

Fixed air . 

Fixed alkali 

Gas f uliginosum 



Present Meaning 

Anhydride (oxide). 

Carbon dioxide. 

Chlorine. 

Carbon dioxide. 

Gas. 

Potassium carbonate. 

Nitric acid. 

Aqua regia. 

Nitrogen. 

Bismuth Subnitrate. 

Oxide. 

Silver carbonate. 

Ferric oxide. 

Oxygen. 

Mercurous chloride. 

Oxygen. 

Carbon dioxide. 

Sodium carbonate. 

Combustible gas. 



APPENDIX E 



387 



Old Terms 

Gas pingue 
Gas siccum 
Gas sylvestre 
Grey calx of lead 
Hartshorn . 
Liver of Sulphur 
Magnesia alba . 
Marcasite . 
Marine acid 
Mephitic air 
Mercurius calcinatus . 
Mercurius dulcis 
Mercurius Niter 
Mercurius precipitatus per 

se . . 
Mercurius precipitatus 

ruber 
Mercurius sublimatus 
Mercurius vitae . 
Mors met alio rum 
Niter . . 
Nitrous air 
Nitrous gas 
Phlogiston . 



Phlogistic air 
Pulvis angelicus . 
Spirit of niter 
Spirit of sulphur 
Spiritus igneo aerius 
Spirit us salis 
Terra pinguis 
Usifur 
Vital air 
Vitriol 

Vitriolated tartar 
Volatile alkali . 



Present Meaning 

Combustible gas. 
Combustible gas. 
Carbon dioxide. 
Lead sesquioxide. 
Ammonia. 

Potassium persulphide. 
Magnesium carbonate. 
Ferric sulphide. 
Hydrochloric acid. 
Nitrogen. 
Mercuric oxide. 
Mercurous chloride. 
Mercuric nitrate. 

Mercuric oxide. 

Mercuric oxide. 

Merc uric, chloride. 

Antimony oxychloride. 

Mercuric chloride. 

Potassium nitrate. 

Nitrogen dioxide. 

Nitrogen dioxide. 

A hypothetical substance, be- 
lieved to exist in all com- 
bustible bodies. 

Nitrogen. 

Antimony oxychloride. 

Nitric acid. 

Sulphuric acid. 

Oxygen. 

Hydrochloric acid. 

Same meaning as phlogiston. 

Artificial mercuric sulphide. 

Oxygen. 

Sulphate. 

Potassium sulphate. 

Ammonium Carbonate. 



388 



MODERN CHEMISTRY 



TABLE OF THE ELEMENTS AND THEIR 
ATOMIC WEIGHTS 



Name 



Aluminum . 
Antimony . 
Argon . . 
Arsenic . . 
Barium . . 
Bismuth 
Boron . . 
Bromine 
Cadmium . 
Caesium . . 
Calcium . . 
Carbon . . 
Cerium . . 
Chlorine 
Chromium . 
Cobalt . . 
Columbium 
Copper . . 
Erbium . . 
Fluorine 
Gadolinium 
Gallium . . 
Germanium 
Glucinum . 
Gold . . . 
Helium . . 
Hydrogen . 
Indium . . 
Iodine . . 
Iridium . . 
Iron . . . 
Krypton . . 



Symbol 



Al 

Sb 

A 

As 

Ba 

Bi 

B 

Br 

Cd 

Cs 

Ca 

C 

Ce 

CI 

Cr 

Co 

Cb 

Cu 

E 

F 

Gd 

Ga 

Ge 

Gl 

Au 

He 

II 

In 

I 

Ir 

Fe 

Kr 



Atomic Weights 



= 16 



27.1 
120. 

39.9 

75. 
137.4 
208.5 

11. 

79.96 
112.4 
133. 

40. 

12. 
140. 

35.45 

52.1 

59. 

94. 

63.6 
166. 

19. 
156. 

70. 

72. 

9.1 
197.2 
4. 

1.01 
114. 
126.85 
193. 

56. 

81.8 



H = l 



26.9 
119.5 
? 

74.45 
136.4 
206.5 
10.9 
79.34 
111.55 
131.9 
39.8 
11.9 
138.0 
35.18 
51.7 
58.55 
93.0 
63.1 
164.7 
18.9 
155.8 
69.5 
71.9 
9.0 
195.7 
? 

1.0 

113. 

125.89 

191.7 

55.5 

? 



APPENDIX E 



389 



TABLE OF THE ELEMENTS AND THEIR ATOMIC 
WEIGHTS— Continued 



Name 



Lanthanum 

Lead . . . . 

Lithium . . . 
Magnesium 

Manganese . . 

Mercury . . . 

Molybdenum . 

Neodymium . 

Neon . . . 

Nickel . . . 

Nitrogen . . 

Osmium . . 

Oxygen . . . 

Palladi um . . 
Phosphorus 

Platinum . . 

Potassium . . 
Praseodymium 

Rhodium . . 

Rubidium . . 
Ruthenium 

Samarium . . 

Scandium . . 

Selenium . . 

Silicon . . . 

Silver . . . 

Sodium . . . 

Strontium . . 

Sulphur . . . 

Tantalum . . 

Tellurium . . 

Terbium . . 



Symbol 


Atomic 


Weights 


= 16 


H = l 


La 


138. 


137.6 


Pb 


206.9 


205.36 


Li 


7. 


6.97 


Mg 


24.36 


24.1 


Mn 


55. 


54.6 


Hg 


200.3 


198.50 


Mo 


96. 


95.3 


Nd 


143.6 


142.5 


Ne 


20. 


? 


Ni 


58.7 


58.25 


N 


14.04 


13.93 


Os 


191. 


189.6 





16. 


15.88 


Pd 


106. 


106.2 


P 


31. 


30.75 


Pt 


194.8 


193.4 


K 


39.15 


38.82 


Pr 


140.5 


139.4 


Rh 


103. 


102.2 


Rb 


85.4 


84.75 


Ru 


101.7 


100.9 


Sm 


150. 


149.2 


Sc 


44.1 


43.8 


Se 


79.1 


78.6 


Si 


28.4 


28.2 


Ag 


107.93 


107.11 


Na 


23.05 


22.88 


Sr 


87.6 


86.95 


S 


32.06 


31.83 


Ta 


183. 


181.5 


Te 


127. 


126.5 


Tr 


16*0. 


158.8 



390 



MODERN CHEMISTRY 



TABLE OF THE ELEMENTS AND THEIR ATOMIC 
WEIGHTS— Continued 





Symbol 


Atomic Weights 




= 16 


H = l 


Thallium 

Thorium 

Thulium 

Tin .... 


Tl 

Th 

Tm 

Sn 

Ti 

W 

u 

V 

X 

Yt 
Y 
Zn 
Zr 


204.1 
232.5 
171. 
118.5 

48.1 
184. 
239.5 

51.2 
128. 
173. 

89. 

65.4 

90.7 


202.61 
230.8 
169.4 
118.1 


Titanium 

Tungsten 

Uranium ....... 

Vanadium 

Xenon 

Ytterbium 

Yttrium 

Zinc 

Zirconium 


47.8 
182.6 
237.8 

51.0 

? 

171.9 

88.3 
64.9 

89.7 



The above table shows two columns of atomic weights; the first as- 
sumes = 16 as the standard, the second, H = 1. 






GLOSSARY OF CHEMICALS AND MINERALS 

Agate. A variety of quartz, occurring often in variegated color 

arranged concentrically. 
alabaster. A fine-grained, white variety of gypsum. 
alum. A double sulphate, of general formula, ]\I 2 R 2 (S0 4 ) 4 24 H 2 0. 
alumina. Aluminum oxide, A1 2 3 . 
amethyst. A variety of quartz, 
anthracite. Natural coal, possessing little or no oil or other volatile 

products. Hard coal. 
antichlor. A reagent used to neutralize chlorine when in excess. 
aragonite. A variety of calcite, CaC0 3 . 
argentite. Native silver sulphide, 
arsenic. The popular name for arsenic trioxide. 
arsenious acid. Another name for arsenic trioxide. 
arsine. Hydrogen arsenide, AsH 3 . 
azurite. An ore of copper, blue in color, composition Cu(OH) 9 , 

2 CuC0 3 . 
Baryta. Barium oxide. 
baryta water. Barium hydroxide, 
bauxite. A hydrated oxide of aluminum, A1 2 3 , H 2 0, used as a source 

for aluminum. 
benzene. A light oil obtained from petroleum, composition C 6 H 6 . 
bicarbonate of soda. Cooking soda, XaHC0 3 . 
bismuth ocher. Bismuth oxide, Bi 2 3 . 
bismuthite. Native bismuth sulphide. 

bituminous. Containing bitumen or oil. Applied to soft coals. 
blanc de fard. Bismuth subnitrate, BiOX0 3 . 
blende. Xative zinc sulphide, 
blue vitriol. Copper sulphate. 
borax. Sodium diborate, Na 2 B 4 7 . 
braunite. Xative Mn 2 3 . 

butter of antimony. An old name for antimony trichloride. 
Calamine. An ore of zinc, Zn 2 Si0 4 , H 2 0. 

391 



392 . MODERN CHEMISTRY 

calchopyrite. A sulphide of iron and copper, Cu 2 S, Fe 2 S 3 . 

calcite. Crystallized calcium carbonate. 

calomel. Mercurous chloride, Hg 2 Cl 2 . 

carbonado. A variety of diamond occurring in black pebbles or 
masses. 

carborundum. A hard substance, made by combining, at high tempera- 
tures, silica and carbon, 
assiterite. Native stannic oxide, Sn0 2 , the chief ore of tin. 
rustic potash. Potassium hydroxide. 

caustic soda. Sodium hydroxide. 

celestite. Strontium sulphate. 

cement. A variety of lime prepared from limestone, containing from 
40 to 50 per cent of slate. 

chalcedony. A variety of quartz. 

chalk. A soft variety of limestone, composed of the shells of diatoms. 

chloride of lime. A common name for bleaching powder. 

chrome alum. A sulphate of potassium and chromium. 

chrome red. Basic lead chromate, Pb 2 Cr0 5 . 

chrome yellow. Lead chromate, PbCr0 4 . 

cinnabar. The chief ore of mercury, HgS. 

clay. A hydrated silicate of aluminum, containing various impurities. 

colcothar. Ferric oxide, Fe 2 3 . 

copperas. Ferrous sulphate. 

corrosive sublimate. Mercuric chloride, HgCl 2 . 

corundum. Anhydrous alumina, uncrystallized. 

cryolite. A fluoride of sodium and aluminum, NaAlF 4 . 

Dolomite. A native carbonate of magnesium and calcium. 

Emerald. (Oriental.) Crystallized alumina, green in color. 

emery. Massive, opaque alumina. 

epsom salts. Magnesium sulphate. 

euchlorine. A solution of chlorine in water. 

Fat lime. Lime made from pure limestone. 

feldspar. A silicate of potassium and aluminum, which, decomposed, 
forms clay. 

fool's gold. Ferric sulphide, Fe 2 S 3 . 

fuller's earth. A variety of clay. 

fuming liquor of Libavius. Anhydrous stannic chloride. 

Galena. The chief ore of lead, PbS. 

green vitriol. Ferrous sulphate. 



GLOSSARY 393 

gypsum. Native calcium sulphate. 

Hartshorn. An old term for ammonia. 

heavy spar. Native barium sulphate. 

hematite. An important ore of iron, of the composition Fe 2 3 . 

horn silver. Native silver chloride. 

hydraulic cement. Lime containing from 10 to 30 per cent of silica, 
having the property of hardening under water. 

hypo. The photographer's name for sodium thiosulphate. 

Iceland spar. A transparent, crystalline variety of calcium carbonate. 

infusorial earth. A grayish white earth, composed largely of silica, 
resulting from the secretion of diatoms. 

Jewelers rouge. An oxide of iron, red in color, used in polishing and 
as a pigment. 

Kaolin. A pure variety of clay, formed by the decomposition of feld- 
spar. 

kelp. The ashes of seaweeds, used as a source of certain potash salts 
and of iodine. 

kerosene. Popularly called coal oil. An oil obtained by the distilla- 
tion of petroleum. 

kieserite. Native magnesium sulphate. 

kupfer nickel. Nickel arsenide, NiAs. 

Labarraque's solution. Sodium hypochlorite. 

lac sulphuris. Sulphur precipitated from a solution of it in lime- 
water. 

laughing gas. Nitrous oxide, N 2 0. 

lean lime. Lime made from impure limestone. 

lime. Calcium oxide, CaO. 

limestone. Calcium carbonate, uncrystallized. 

lime-water. Calcium hydroxide. 

litharge. Impure lead oxide, PbO. 

lunar caustic. A commercial term for silver nitrate. 

Magnesia. Magnesium oxide. 

magnesite. Native magnesium carbonate. 

magnetic pyrites. A mixture of FeS and Fe 2 S 3 . This mixture is 
given its name because of magnetic properties. 

malachite. An ore of copper, CuC0 3 , Cu(OH) 2 . 

marble. Crystallized limestone. 

marcasite. A variety of ferric sulphide, FeS 2 . 

massicot. Lead oxide, PbO. 



394 MODERN CHEMISTRY 

milk of lime. Calcium hydroxide, containing more or less lime in 

suspension. 
milk of sulphur. Same as lac sulphuris. 
minium. Red lead, Pb 3 4 . 

mispickeL An important ore of arsenic, FeSAs, 
Naphtha. A light oil, obtained from petroleum. 
Nessler's solution. A solution used in testing for ammonia, 
niter. Another name for potassium nitrate. 
Nordhausen's acid. The same as fuming sulphuric acid, H 2 S 2 7 . 
Oil of vitriol. Sulphuric acid. 
opal. A variety of silica, Si0 2 . 
oriental. A term applied to the true emerald and certain other gems, to 

distinguish them from less valuable stones similar in appearance. 
orpiment. A sulphide of arsenic, yellow in color, having composition 

As 2 S 3 . 
Paraffin. A wax obtained in the later distillation of petroleum. 
Paris green. A popular name for Scheele's and Schweinfurth's green, 

compounds of arsenic. 
pearl ash. Pure potassium carbonate. 
pearl white. Bismuth oxychloride, BiOCl. 

petroleum. Rock oil, found native in various parts of the world, 
plaster of Paris. Calcined calcium sulphate. 
plastic sulphur. A dark-colored, allotropic form of sulphur, somewhat 

resembling rubber. 
potash. Another name for commercial potassium carbonate. Also a 

loose name for potassium chlorate. 
powder of Algaroth. A variable compound of antimony, approximately 

SbOCl. 
purple of Cassius. A purplish-colored precipitate obtained in testing 

a solution of gold with stannous chloride. 
pyrites. A common name for ferric sulphide, FeS 2 . 
pyrolusite. Native manganese dioxide. 
Quartz. Silicon dioxide. 
quicklime. The same as lime. 
Realgar. Red sulphide of arsenic, As 2 S 2 . 
red lead. The same as minium. 
red precipitate. Mercuric oxide. 

rose quartz. A variety of quartz, somewhat pink in color. 
Sal ammoniac. Ammonium chloride. 






GLOSSARY 395 

sal soda. Sodium carbonate. 

salt. The common name for sodium chloride. 

salt cake. Sodium sulphate. 

saltpeter. Potassium nitrate. 

sapphire. Crystallized alumina. 

Scheele's green. Copper arsenite, CuHAs0 3 . 

silica. Silicon dioxide. 

slaked lime. Lime treated with water. 

smalt. A silicate of cobalt aud potassium. 

smoky quartz. A variety of silica, brown or smoky in color. 

soda. Same as sal soda. 

soda, cooking. Same as sodium bicarbonate, NaHC0 3 . 

spathic iron. Native iron carbonate, FeC0 3 . 

specular iron. A variety of hematite. 

spiegeleisen. A variety of iron containing manganese and carbon. 

stibine. Same as antimoniureted hydrogen, SbH 3 . 

strontianite. Native strontium carbonate. 

subnitrate of bismuth. Basic bismuth nitrate, BiON0 3 . 

sugar of lead. Lead acetate. 

Topaz. Crystallized alumina with small quantity of coloring matter. 

Vermilion. Artificial mercuric sulphide. 

White arsenic. Arsenic trioxide. 

white lead. Basic lead carbonate, used as a paint. 

white vitriol. Zinc sulphate. 

witherite.- Native barium carbonate. 

Zinc white. Zinc oxide, ZnO, used as a paint. 

GLOSSARY OF TECHNICAL TERMS IN CHEMISTRY 

Acidify. To make acid. 

acidulate. To add acid to, until no longer alkaline or neutral. 

actinic. Referring to light rays, having the power to effect chemical 

changes. 
air-bath. A small oven used for drying substances. 
alkali. A compound of hydrogen, oxygen, and some metallic element, 

soluble in water, having the power to neutralize acids ; as caustic 

soda, NaOH. 
allotropic. Literally, another form; a term applied to the unusual 

form of an element. 



396 MODERN CHEMISTRY 

allotropism. The phenomenon of existing in two or more forms, 
alloy. The product resulting from fusing together two or more metals, 
amalgam. An alloy, one constituent of which is mercury, 
amorphous. Without any special form, uncrystallized, massive, 
anaesthetic. An agent used to produce insensibility. 
anhydride. An oxide, usually non-metallic, which forms some acid 

upon the addition of water, 
anhydrous. Without water. An anhydrous salt is one from which 

the water of crystallization has been removed, 
anion. A negative ion. See ion. 
antiseptic. A substance used to prevent decay, or to destroy noxious 

germs, 
argentiferous. Silver-bearing. 
aspirator. Apparatus used to secure the passage of air or any other 

gas through certain vessels. 
assay. Determination of the quantity of the various constituents of 

a metallic ore. 
Basic. Having the properties of an alkali or base, 
binary. A compound consisting of two elements, 
brightening. The sudden brilliant appearance of the silver assay when 

the lead has all been removed by cupellation. 
bumping. A term applied to the violent boiling of the liquid in a 

vessel, causing it to jump, 
burette. A graduated tube, with stop-cock, used in volumetric work 

for measuring accurately a liquid. 
Calcine. To heat strongly, 
carbureting. Adding hydrocarbon compounds to an illuminating gas, 

as in making water gas. 
cathion. An electropositive ion. 
cementation. An old process of making steel by imbedding wrought 

iron in powdered charcoal and heating several days, 
chemism. The so-called affinity that one element or substance has 

for another, 
commercial. A term applied to chemicals as usually furnished to the 

trade; not absolutely pure; in distinction from chemically pure 

reagents. 
concentrated. Strong; undiluted, 
converter. A large, egg-shaped furnace, used in making steel from 

cast iron and in purifying copper. 



GLOSSARY 397 

c. p. Chemically pure. 

crucible. A small vessel, made to withstand great heat. Named 
from the Latin word crux, because the old alchemists thus marked 
their crucibles. 

crystalline. Composed of crystals. 

cupel. A small cup, made of bone ashes ; used by assayers in deter- 
mining the gold and silver in an ore. 

cupellation. The process of separating lead and silver by the oxida- 
tion of the former. 

Decant. To pour off the liquid from a precipitate, after the latter 
has settled. 

decrepitate. To burst in pieces with a crackling sound, as many salts 
do when heated with the blowpipe. 

deflagrate. To burn vigorously. 

deflagrating spoon. A small metallic cup or spoon with a long wire 
handle attached. Used for holding combustible substances when 
burning in oxygen or other gases. 

deliquesce. To take up moisture from the air. 

deoxidizing agent. See reducing agent. 

desiccate. To dry. 

desiccator. A vessel used in drying or keeping dry a substance which 
is to be weighed accurately. 

destructive distillation. The process of heating in closed retorts a 
substance to such a temperature as to effect its decomposition. 

digest. To warm gently. 

disinfectant. A substance used to cleanse and purify unwholesome 
places, as well as to destroy disease germs. 

displacement. A method of collecting a gas in a vessel filled with 
air, or some other gas, depending upon the difference in density 
of the two. 

dissociate. To break up a compound body into parts. 

distill. To evaporate a liquid and condense again in another vessel. 

distillate. The liquid obtained in the process of distillation. 

dyad. An element having a valence of two. 

Ebullition. Rapid boiling. 

effervescence. The act of bubbling, as seen upon the application of 
an acid to a carbonate. 

effloresce. To give up at ordinary temperatures the water of crystal- 
lization. 



398 MODERN CHEMISTRY 

electrode. The terminal of a battery. 

electro-positive. A term applied to elements attracted to the negative 

electrode, 
equivalence. A term sometimes used instead of valence. 
escharotic. An agent which corrodes or destroys ; a caustic. 
evolve. To set free. 
excess. A quantity more than sufficient to secure certain chemical 

action. 
Filtrate. The liquid obtained after passing through the filter paper, 

in removing the precipitate. 
fixed. The opposite of volatile. 
flocculent. Flaky, 
flux. Any substance used to lower the melting point of another ; as 

limestone with iron ore in the blast furnace. 
formula. A combination of symbols used to represent a molecule of 

a compound body. 
fractional distillation. The process of separating by distillation the 

several constituents of a mixture of liquids, by means of their 

different boiling points. 
Gangue. The impurities contained in an ore or mineral, 
gelatinous. Like starch paste in appearance. 
generate. To produce or set free, as a gas. 
germicide. A substance used to destroy bacteria or germs. 
granulated. In irregularly shaped small particles, secured by pouring 

the fused metal into cold water, 
graphitoidal. Resembling graphite. 
gravimetric. Measurement or estimation by weight. 
Halogen. Literally, salt producer; applied to the members of the 

chlorine group. 
hydrated. Containing water. 
hydroxyl. A term applied to the radical OH. 
hygroscopic. Applied to substances which readily absorb moisture 

from the air. 
Ignite. To set fire to. 
indicator. A substance used to show the completion of a chemical 

reaction. 
inflammable. Combustible. 
ion. An atom or group of atoms in a solution, which serves as a 

carrier of electricity. 






GLOSSARY 399 

ionization. The separation of a substance into ions. 

isomeric. Applied to substances having the same percentage composi- 
tion, though differing in characteristics. 

isomorphous. Of the same crystalline form. 

Leach. To treat with water ; to remove the soluble salts from a 
mixture of substances by means of water. 

liquation. The process of separating one metal from another by 
cautiously fusing, so that one will flow out before the melting 
point of the other is reached. 

lixiviate. Synonymous with leach. 

lute. To seal air-tight. 

Manipulation. Setting up or arranging apparatus for experiment. 

matte. A mixture of metallic sulphides obtained in the early 
stages of the reduction of copper ores, containing lead, silver, 
etc. 

meniscus. The upper curved surface of a liquid contained in a small 
tube. 

monad. An element the valence of which is one. 

mono-basic. A term applied to an acid having only one replaceable 
atom of hydrogen. 

mordant. A substance used to set the color in dyeing. 

mother liquor. The liquid remaining after the principal salt con- 
tained in solution has been removed by crystallization. 

Nascent. Applied to a gas when first liberated from its compound. 
It is believed to exist then in the atomic condition. 

native. Not in combination, free. 

neutral. Neither acid nor alkaline. 

neutralization. The combination of an acid with an alkali so as to 
destroy the properties of each, and produce a salt. 

nitrogenous. Containing nitrogen. Organic matter containing nitro- 
gen is thus characterized. 

Occlude. To condense upon the surface or within the pores. Especially 
seen in the action of platinum upon hydrogen. 

oxidation. The union of a substance with oxygen. 

oxidizing agent. A substance which readily gives up a portion of its 
oxygen to combine with some other substance. 

oxygenized. Containing considerable oxygen. 

Paste. A special variety of glass, used sometimes for making imita- 
tion diamonds. 



400 MODERN CHEMISTRY 

pigs. The term applied to cast iron as molded when first drawn from 
the blast furnace. Applied also to the molds themselves. 

pipette. A small graduated glass tube used in measuring small 
quantities of a liquid. 

pneumatic. Pertaining to gases; applied to the trough or pan used 
in collecting gases. 

polymerism. A term referring to the cases of compounds which have 
the same percentage composition, but different molecular weights. 

precipitate. A solid thrown down in a liquid by some reagent. 

Qualitative analysis. The determination of the kind of matter which 
enters into a substance. 

quantitative analysis. The determination of the amount of a sub- 
stance contained in a compound. 

Radical. A group of elements which seems to act as a single ele- 
ment. 

reaction. The action of two or more substances upon each other. 

reagent. A substance used to bring about some chemical change. 

reducing agent. A substance used to convert a compound from a 
higher to a lower order, as from an ic to an ous compound ; or, 
to remove the oxygen from an oxide. 

residual. That which remains. 

reverberatory. A variety of furnace, usually of low, arching ceiling. 
See Fig. 57 in text. 

roast. To heat strongly; to oxidize metallic ores, expelling the 
sulphur as S0 2 . 

Sand-bath. A small iron saucer containing sand, used the same as a 
wire screen in protecting glassware when being heated. 

saturated. Full} 7 satisfied; containing all it can hold. 

scintillate. To burn with sparks. 

siliceous earth. Material consisting largely of silica. 

slag. The dark-colored glass formed in the reduction of metallic ores 
from the flux used and the gangue present. 

solvent. A liquid which dissolves some particular substance. 

spit. Silver on being melted absorbs considerable oxygen. Upon 
cooling it again expels this, sometimes with considerable energy, 
throwing out fine particles of the molten silver. This is termed 
spitting. 

stable. Not easily decomposed. 

sublimate. The substance obtained by sublimation. 



GLOSSARY 401 

sublimation. The vaporizing of a solid and recondensing. The same 
in reference to solids that distillation is with liquids. 

supernatant. Said of a liquid overlying a precipitate after the latter 
has subsided. 

suspension. Said of a solid in the form of fine particles floating 
throughout the liquid. 

symbol. A letter or letters representing an atom of an element. 

Thio. From a Greek word, meaning sulphur. 

treat. To apply or add to. 

triad. An element having a valence of three. 

tubulated. Applied to a flask having a small tube-like opening in the 
side, fitted with a stop-cock. 

tubulure. A small, tube-like opening. 

tuydre. A blast or air pipe for conducting the strong currents of air 
into the blast furnace. 

Valence. The power which an atom or group of elements has of com- 
bining with some other element taken as a standard. 

volatile. Easy to vaporize. 

volatilize. To drive off in the form of vapor. 

volumetric. Estimation of the quantity of a substance by measuring. 



'*! 



< 



INDEX 



Absolute thermometer, 95. 
Absolute zero, 95. 
Acetic acid, test for, 345. 
Acetylene, burners for, 150. 

characteristics of, 150. 

experiments with, 151, 152. 

generators, 149. 

preparation of, 148. 
Acids, 125. 

classes of, 343. 

composition of, 126. 

detection of, 343. 

nomenclature of, 128. 

preliminaries to testing, 345, 347. 

properties of, 125. 
Air, estimation of its constituents, 
350, 

estimation of weight, 97. 

liquefaction of, 97. 
Air-slaked lime, 222. 
Alchemistic terms, 386. 
Alkali earths, 219. 
Alkalies, 125. 
Alkali metals, 207. 
Allotropism. 59. 
Aluminum, 263. 

bronze, 235. 

characteristics of, 264. 

hydroxide, 269. 

source of supply, 263. 

test for, 337. 

uses, 264. 
Alums, 266. 

kinds of, 267. 

preparation of, 266. 

uses of, 267. 

uses of, for clarifying water, 268. 



Amalgams, 258. 

methods of making, 258. 
Ammonia, 73. 

absorption of, by charcoal, 78. 

as a refrigerant, 78. 

characteristics of, 76. 

commercial supply, 74. 

decomposition of, by platinum, 78. 

estimation of the constituents, 351. 

fountain, 77. 

preparation for commerce, 74. 

test for, 342. 

uses, 78. 
Ammonium. 67. 
Anhydride, 83. 
Anions, 329. 
Antichlor, 112. 

Antimoniureted hydrogen, 292. 
Antimony, 290. 

amorphous, 292. 

black, 292. 

characteristics of, 291. 

chloride, 293. 

oxides, 293. 

oxychloride, 293. 

reduction of ore, 290. 

sulphide, 294. 

test for. 334. 

uses, 292. 
Apparatus for pupils, 357. 
Aqua regia, 88. 
Argentite, 238. 
Argon, characteristics of, 90. 

discovery, 89. 
Arrangement of bottles, 356. 
Arsenic, characteristics of, 286. 

Marsh's test for, 287. 



403 



404 



INDEX 



oxides, 288. 

reduction of ores, 285. 

source of supply, 285. 

sulphide, 290. 

uses of, 286. 
Arsenical pyrite, 285. 
Arsine, 287. 
Asbestos, 219. 
Aspirators, 370. 
Atmosphere, 91. 
Atom, definition of, 11. 
Atomic weights, 68. 

determination of, 198. 
Avogadro's Law, 196. 

application of, 198. 

proof of, 196. 
Azurite, 233. 

Banca tin, 270. 
Barium, 229. 

carbonate, 229. 

chloride, 229. 

hydroxide, 230. 

nitrate, 229. 

separation from calcium, 340. 

sulphate, 229. 

tests for, 340. 
Barometric reading, correction of, 

374. 
Baryta, 229. 
Base, 124. 
Bauxite, 264. 
Bell metal, 235. 
Bessemer steel, 303. 
Binary compounds, 131. 
Biographical appendix, 382. 
Bismuth, 294. 

characteristics of, 294. 

compounds, classes of, 295. 

nitrate, 295. 

ocher, 296. 

oxychloride, 296. 

trichloride, 296. 

trioxide, 296. 

uses, 295. 



Bismuthite, 294. 

Bismuthyl compounds, 295. 

Black ash, 211. 

Black lead, 136. 

Blast furnace, 300. 

Bleaching powder, 227. 

Bloom, 302. 

Blowpipe work, 361. 

Blue prints, 245. 

Blue vitriol, 235. 

Bohemian glass, 188. 

Bordeaux mixture, 236. 

Bornite, 233. 

Bottles, opening of, 366. 

Boyle's Law, 93. 

Brass, 235. 

Bromides, test for, 344. 

Bromine, characteristics of, 117. 

commercial supply, 116. 

experiments with, 118. 

occurrence of, 116. 

preparation of, 117. 

test for, 117. . 
'uses, 118. 
Bronze, 235. 
Burnt alum, 267. 

Cadmium, 254. 

characteristics of, 255. 

nitrate, 256. 

reduction of, 255. 

sulphide, 256. 

test for, 334. 
Calchopyrite, 233. 
Calcite, 221. 
Calcium, 220. 

carbide, 148. 

carbonate, 223. 

characteristics of, 221. 

chloride, 224. 

history of, 221. 

hydroxide, 223. 

oxide, 221. 

sulphate, 224. 
Calomel, 260. 



INDEX 



405 



Carbon, abundance of, 135. 

as an absorbent, 139. 

as a reducing agent, 139. 

forms of, 135. 

uses of, 140. 
Carbon dioxide, 142. 

characteristics of, 144. 

estimation of, 348. 

liquid, 144. 

preparation of, 143. 

source of, 142. 

uses of, 144. 
Carbon monoxide, 141. 
Carre's ice machine, 79. 
Cassiterite, 270. 
Cast iron, 302. 

Castner's process for sodium, 208. 
Catalysis, 51. 
Cathions, 329. 
Caustic soda, 209. 
Cements, 225. 

composition of, 226. 
Chamber acid, 181. 
Charcoal, 137. 
Charles's Law, 94. 

problems with, 96. 
Chemical changes, 15. 

experiments to illustrate, 15, 16, 
17, 18. 
Chloric acid, 345. 
Chlorine, as a bleaching age tit, 111. 

characteristics of, 109. 

chemistry of its preparation, 106. 

Deacon's process, 106. 

experiments with, 108. 

history of, 102. 

liquid, 110. 

occurrence, 103. 

preparation, 103. 

uses of, 111. 

water, 105. 

Weldon's process, 104. 
Choke damp, 144. 
Chromic acid, 321. 
Chromium, 317, 



compounds, 317. 

conversion of compounds, 319. 

hydroxide, 321. 

oxides, 320. 

test for, 337. 

uses of, 321. 
Chromite, 317. 
Cinnabar, 257, 260. 
Clay, 265. 
Coal, 137. 
Coal gas, 153. 
Cobalt, 311. 

compounds, 311. 

test for, 338. 
Coke, 138. 

Combination, laws of, 166. 
Combining weights, 164. 
Combustible substances, 57. 
Combustion, 56. 
Compounds, 10. 

saturated, 24. 
Converter, 303. 
Copper, 232. 

alloys of, 235. 

blister, 233. 

characteristics of, 234. 

pyrite, 233. 

reduction of, 233. 

salts, 235. 

supply of, 232. 

tests for, 233. 
Copperas, 308. 
Corals, 221. 

Corrosive sublimate, 260. 
Corundum, 265. 
Crocosite, 317. 
Crown glass, 188. 
Cryolite, 264. 
Cupel, 239. 
Cupellation, 239. 
Cupola furnace, 304. 
Cupric acetylide, 236. 

chloride, 236. 

nitrate, 236. 

oxide, 237, 



406 



INDEX 



sulphate, 235. 
sulphide, 236. 
Cyanide process for gold, 247. 

Decanting, 364. 

Definite Proportions, Law of, 158. 

Deliquescent bodies, 31. 

Delivery tubes, preparation of, 358. 

Dewar bulbs, 97. 

Diamonds, 135. 

practical uses, 136. 
Diatomic molecules, 200. 
Diffusion of gases, 92. 
Dissociation, 329. 
Distillation, destructive, 137. 

fractional, 137. 
Dolomite, 219. 

Downward displacement, 362. 
Drying of tubes, 375. 
Dyads, 23. 
Dynamite, 89. 

Efflorescent bodies, 30. 
Electrolysis of water, 32. 
Electrolytic apparatus, 367. 
Elements, classes of, 204. 

definition of, 8. 

table of, 8, 204, 388. 

vacancies in table, 206. 

valence of, 8. 
Emerald, 265. 
Emery, 265. 
Epsom salts, 220^_ _ 
Equations, 27. 

exercise in, 28. 

writing, 69. 

value of, 69. 
Etching glass, 102. 
Ethane, 147. 
Euchlorine, 105. 
Eudiometer, 33, 369. 
Experiments, recording, 376. 

Feldspar, 265. 
Ferric chloride, 308. 



oxide, 309. 

salts, how changed to ferrous, 307. 

salts, how distinguished, 306. 

sulphate, 308. 

sulphide, 308. 
Ferrous salts, how changed to ferric, 
307. 

how tested, 306. 
Fertilizers, 194. 
Filter flask, 370. 
Filtering, 364. 
Fire damp, 146. 
Fixing bath, 244. 
Flame, 56. 

Flame tests for barium, etc., 230. 
Flint glass, 188. 
Fluorine, 101. 

compounds of, 102. 
FooFs gold, 300. 
Formulae, determination of, 201. 

meaning of, 66. 
Franklinite, 250. 

Galena, 274. 
Ganister, 303. 
Gas carbon, 138. 
Gas generators, 373. 
Gases, collecting, 362. 

illuminating, 152. 
German silver, 253. 
Glacial phosphoric acid, 194. 
Glass, 187. 

annealing, 189. 

cutting, 358. 

etching, 102. 

manufacture of, 187. 

varieties of, 188. 
Glauber's salt, 212. 
Glossary of chemicals and min- 
erals, 391. 
Glossary of technical terms, 395. 
Gold, 246. 

characteristics of, 248. 

methods of mining, 246. 
Graphite, 136. 



^- 



INDEX 



407 



Greek fire, 175. 
Green fire, 229. 
Greenockite, 255. 
Green vitriol, 308. 
Guncotton, 89. 
Gunpowder, 175. 

separation of, 19. 
Gypsum, 221. 

Halogens, 101. 

comparison of, 122. 
Hard waters, 226. 
Harveyized steel, 310. 
Heavy spar, 229. 
Hematite, 299. 
Horn silver, 238. 
Hydraulic cement, 225. 
Hydraulic mining, 246. 
Hydriodic acid, test, 344. 
Hydrobromic acid, test, 344. 
Hydrocarbons, 146. 
Hydrochloric acid, characteristics 
of, 115. 

commercial supply, 113. 

composition of, proof, 354. 

composition of, estimation, 352. 

experiments with, 114. 

history of, 112. 

preparation of, 112. 

test for, S44. 

uses, 115. 
Hydrogen, 36. 

characteristics of, 42. 

experiments with, 42. 

liquid, 45. 

methods of preparing, 36. 

occlusion of, 44, 314. 

uses, 45. 
Hydrogen dioxide, 63. 
Hydrogen sulphide, 175. 
Hydroxides, 125. 
Hydroxyl, 66. 

Iceland spar, 221. 
Ice machine, 78, 80. 



Ice manufacture, 79. 
Illuminating gases, 152. 

composition of, 155. 

manufacture of, 153. 
Iodides, test for, 344. 
Iodine, 119. 

characteristics of, 12 1. 

experiments with, 121. 

preparation of, 119. 

solvents for, 122. 

uses, 122. 
Ionic theory, 328. 
Ionization, 329. 
Ions, 329. 
Iridium, 314. 
Iron, cast, 302. 

compounds of, 305. 

distribution of, 299. 

forms of, 305. 

protoxide, 309. 

pyrite, 300. 

reduction of, 302. 

test for, 337. 

uses, 305. 

wrought, 302. 

Jets, preparation of, 359. 

Kaolin, 265. 

Kindling temperature, 57. 

Laboratory suggestions, 355. 
Lampblack, 139. 
Lead, 274. 

acetate, 279. 

carbonate, 280. 

characteristics of, 277. 

chloride, 279. 

chromate, 282. 

nitrate, 279. 

oxides, 280. 

reduction of ores, 275. 

sulphate, 279. 

sulphide, 282. 

tests for, 283, 331, 

uses, 277. 



408 



INDEX 



Leblanc's process for soda, 211. 
Lime, 221. 

properties of, 222. 
Limonite, 300. 
Linde's apparatus, 98. 
Liquid air, 100. 

apparatus for, 98. 
Liter, weight of, 201. 
Litharge, 280. 
Lithium, test for, 342. 
Lunar caustic, 243. 

Magnesia, 220. 
Magnesium, 219. 

compounds, 220. 

test for, 340. 
Magnetite, 299. 
Malachite, 233. 
Manganese, 323. 

compounds of, 323. 

dioxide, 324. 

dioxide, as a catalytic agent, 
50. 

test for, 338. 
Manganic acid, 324. 
Marsh gas, 146. 
Marsh's test for arsenic, 287. 
Matte, 233. 
Matter, 11. 

theories of, 8. 
Measurements, 363. 
Meerschaum, 219. 
Mercuric chloride, 260. 

nitrate, 259. 

oxide, 260. 

salts, distinguished, 261. 

sulphide, 260. 
Mercurous chloride, 260. 

nitrate, 259. 
Mercury, 256. 

characteristics of, 257. 

reduction of, 257. 

solvents for, 259. 

tests for, 331, 333. 

uses, 259, 



Metals, 203. 

displacing power of, 169. 
Metaphosphoric acid, 194. 
Meteorites, 299. 
Methane, 146. 
Micro-crith, 68. 
Minium, 280. 
Mixtures, 18. 
Molecular weights, 69. 
Molecules, 11. 

constitution of, 199. 

of compound bodies, 12. 
Monads, 23. 

Monatomic molecules, 200. 
Monobasic acids, 194. 
Mortar, 222. 

Multiple Proportions, Law of, 
161. 

Natural gas, 155. 
Negatives, photographic, 244. 
Neutralization, 124. 
Nickel, 309. 

compounds of, 310. 

tests for, 311, 338. 

uses, 310. 
Nitric acid, characteristics of, 87. 

in the air, 86. 

preparation of, 86. 

test for, 85. 

uses, 88. 
Nitric oxide, 82. 

characteristics of, 83. 
Nitrogen, 71. 

characteristics of, 73. 

oxides of, 81. 

pentoxide, 86. 

tetroxide, 85. 

uses of, 73. 
Nitrogen group, 297. 
Nitroglycerine, 89. 
Nitrous acid, preparation, 84, 

test for, 84. 
Nitrous anhydride, 83. 

oxide, 81, 



INDEX 



409 



Occlusion, 44, 314. 
Oil of vitriol, 179. 
Olefiant gas, 147. 
Orpiment, 285. 
Osmium. 314. 
Oxidation, 56. 
Oxides, 132. 
Oxidizing flame, 361. 
Oxygen, 47. 

characteristics of, 54. 

determination of weight, 55. 

experiments with, 49. 

liquid, 54. 

Motay's method, 52. 

preparation, 48. 

preparation from potassium per- 
manganate, 53. 

uses, 55. 
Oxy-hydrogen blowpipe, 58. 
Ozone, 59. 

characteristics of, 61. 

comparison with oxygen, 60. 

liquid, 61. 

Palladium. 314. 

Panning gold, 246. 

Paris green. 289. 

Parke's method, 239. 

Paste, 188. 

Pattison's method, 239. 

Pearl ash, 216. 

Pearl white, 297. 

Pentads, 23. 

Periodic Law, 204. 

Phosphates, 194. 

Phosphine, 192. 

Phosphoric acid, test for, 344. 

Phosphorus, 190. 

acids of, 194. 

characteristics of, 191. 

manufacture of, 190. 

oxides of, 193. 

uses of, 192. 
Photographic papers, 244, 

plates, 24?. 



Photography, 243. 
Physical changes, 12. 

illustration of, 13. 

experiments, 12, 14. 
Pig iron. 302. 
Pintsch gas, 154. 
Placer mining, 246. 
Plaster of Paris, 224. 
Platinum, 314. 

alloys of, 315. 

spongy, 314. 

uses, 316. 

wires, 366. 
Polymerism, 62. 
Porcelain, 265. 
Potassium, 214. 

bromide, 217. 

carbonate, 216. 

chlorate, 216. 

chr ornate, 318. 

dichromate, 318. 

hydroxide, 215. 

iodide, 217. 

nitrate, 217. 

permanganate, 324, 325. 

tests for. 217, 342. 
Precipitates, 364. 
Prefixes, per, pro, etc., 133. 
Pressure, standard, 93. 
Puddling furnace, 302. 
Pyrite, 300. 
Pyrolusite, 323. 

Qualitative analysis, 327. 
Quantitative experiments, carbon 

dioxide, estimation of, 348. 
combining weight of copper, 

164. 
combining weight of tin, 165. 
composition of air, 350. 
composition of ammonia, 351. 
composition of hydrochloric acid, 

352. 
definite proportions, proof of law, 

158, 159, 160, WX. 



410 



INDEX 



displacing power,^.. aluminum, 
170. ^ 

magnesium, j?71. 
zinc, 170. 

electrolysis of water, 32. 

manganese dioxide as a catalytic 
agent, proof. of, 51. 

multiple proportions, proof of 
law, 162. 

strength of acid, determination of, 
167. 

strength of alkali, determination 
of, 167. 

strength of salt solution, determi- 
nation of, 169. 

synthesis of water, 33, 34. 

water of crystallization, determi- 
nation of, 349. 

weight of 1 liter of air, 97. 

weight of 1 liter of oxygen, 55. 
Quicksilver, 256. 

Radicals, 66. 
Reactions, 67. 
Realgar, 285. 
Red fire, 228. 
Red precipitate, 260. 
Reducing flame, 361. 
Reference library, 381. 
Rose quartz, 186. 
Ruby, 265. 

Safety lamp, 147. 
Saltpeter, 217. 

Chile, 212. 
Salts, acid, 127. 

definition of, 127. 

formulae of, 128. 

nomenclature of, 130. 

normal, 127. 
Sapphire, 265. 
Scheele's green, 289. 
Scheele's test for arsenic, 290. 
Separation of metals, arsenic, an- 
timony, tin, 334. 



barium, strontium, calcium, mag- 
nesium, 340. 

bismuth, copper, cadmium, 333. 

iron, chromium, aluminum, 337. 

lead, silver, mercury, 330. 

nickel, cobalt, manganese, zinc, 338. 

sodium, potassium, lithium, 342. 
Shot, 278. 
Siderite, 300. 
Silica, 186. 
Silicates, 186. 
Silicic acid, 187. 
Silicon, 184. 
Silver, 238. . 

characteristics of, 241. 

chloride, 242. 

chromate, 242. 

experiments with, 240. 

nitrate, 241. 

Parke's process, 239. 

Pattison's process, 239. 

reduction of ores, 238. 

test for, 332. 

uses, 241. 
Smalt, 312. 
Smithsonite, 250. 
Smoky quartz, 186. 
Soap, 212. 

hard and soft, 213. 
Sodium, 207. 

carbonate, 210. 

characteristics of, 208. 

chloride, 209. 

effects upon water, 38. 

experiments with, 38. 

hydroxide, 209. 

nitrate, 212. 

preparation of, 210. 

sulphate, 212. 

tests for, 214, 342. 
Sodors, 145. 
Solder, 278. 
Solubility of salts, 346. 
Solutions, preparation of, 376. 
Solvay process for soda, 210. 



INDEX 



411 



Sparklets, 145. 
Spathic iron, 300. 
Spelter, 253. 
Spiegeleisen, 304. 
Stalactite, 231. 
Stannic chloride, 272. 

oxide, 274. 

sulphide, 273. 
Stannous chloride, 272. 

sulphide, 273. 
Steel. 303. 

basic lining process, 304. 

comparison with cast iron, 305. 

tempering. 305. 
Stibine, 292. 
Stibnite, 290. 
Strass. 188. 
Strontianite, 228. 
Strontium, 228. 

carbonate, 228. 

hydroxide, 229. 

nitrate. 228. 

separation of, 340. 

tests for, 340. 
Sugar of lead, 279. 
Suint, 214. 

Sulphides, test for, 304. 
Sulphur, 171. 

acids of, 179. 

allotropic, 174. 

characteristics of, 173. 

forms of, 174. 

oxides of, 177. 

source of supply, 172. 

uses, 175. 
Sulphur dioxide, 177. 

characteristics of, 178. 

uses. 179. 
Sulphuric acid, 179. 

characteristics of, 182. 

manufacture of. 181. 

test for, 180, 343. 

uses, 183. 
Sulphurous acid, 183. 

test for, 344. 



Supplies needed, 378. 
Supporters of combustion, 57. 
Symbols, 65. 
Sympathetic inks, 312. 
Synthesis of water, 33. 

Tables : 

comparison of oxygen and ozone, 
62. 

comparison of metals and non- 
metals, 203. 

composition of cements, 226. 

compounds of chromium, 318. 

compounds of manganese, 326. 

iron salts, distinctions between,306. 

names of elements, 9, 388. 

nitrogen group, 297. 

periodic law, 204. 

salts of mercury, 261. 

separation of metals, 

group I, 331. 
group II, 336. 
group III, 339. 
group IV, 341. 
group V, 343. 

tension of aqueous vapor, 374. 

three forms of iron, 305. 

valence, 27. 
Ternary compounds, 131. 
Test-tube repairing, 360. 
Tetrads, 23. 
Thiosulphuric acid, 184. 

test for, 344. 
Tin, 270. 

alloys of, 272. 

characteristics of, 270. 

foil, 272. 

plate, 272. 

salts, 272. 

stone, 270. 

test for, 335. 

uses, 272. 
Triads, 23. 
Tribasic, 194. 
Type-metal, 278. 



412 



INDEX 



Univalent atoms, 23. 
Upward displacement, 362. 

Valence, definition of, 21. 

double, 25. 

exercise in, 26. 

illustration of, 22. 

of radicals, 25. 

variation of, 23. 
Vapor density, determination of, 

200. 
Vein mining for gold, 246. 
Vermilion, 260. 
Vitriol, oil of, 179. 

blue, 335. 

green, 308. 

white, 253. 

Wash bottle, preparation of, 359. 
Water, abundance of, 29. 

analysis of, 32, 34. 

characteristics of, 31. 

composition of, 32. 

decomposition by sodium, 37. 

forms of, 29. 

solvent powers of, 32. 

synthesis of, 33. 
Water gas, 154. 
Water glass, 187. 



Water of crystallization, 30. 

estimation of, 349. 

proof of, 30. 
Weldon's mud, 105. 
White arsenic, 288. 
White lead, 280. 

Dutch method of preparation, 280. 

electrolytic method, 281. 

Milner's method, 281. 
White vitriol, 253. 
Witherite, 229. 
Wrought iron, 302. 

Zinc, 250. 
alloys of, 253. 
blende, 250. 
characteristics of, 252. 
chloride, 253. 
hydroxide, 254. 
ores, 250. 
oxide, 254. 

reaction with acids, 40. 
reduction of, 250. 
sulphate, 253. 
sulphide, 254. 
test for, 338. 
uses, 253. 
white, 254. 



UL [1382 



